Oxidation  and  Reduction  in  Organic  Chem- 
istry from  the  Standpoint  of 
Potential  Differences 

The  System  Hydroquinone  and  Quinone 


By 
FREDERIC  STEARNS  GRANGER,  Ph.D. 


COLUMBIA  UNIVERSITY  PRESS 

1920 
All  rights  reserved 


Oxidation  and  Reduction  in  Organic  Chem- 
istry from  the  Standpoint  of 

Potential  Differences 

i 

The  System  Hydroquinone  and  Quinone 


By 
FREDERIC  STEARNS  GRANGER,  PhJX 


fork 

COLUMBIA  UNIVERSITY  PRESS 

1920 
All  rights  reserved 


Copyright,  1920 
By  COLUMBIA  UNIVERSITY  PRESS 


Printed  from  type,  March  1920 


TO  MY  MOTHER 


435324 


ACKNOWLEDGMENT 

The  author  wishes  to  express  to  Prof.  John  M.  Nelson  his  apprecia- 
tion of  his  kindly  encouragement  and  valued  advice  and  aid  in  the  carry- 
ing out  of  this  work. 


ABSTRACT  OF  DISSERTATION. 

(1)  What  was  attempted? 

(2)  In  how  far  were  the  attempts  successful? 

(3)  What  contribution  actually  new  to  the  science  of  Chemistry  has 
been  made? 

(1)  It  has  been  undertaken  (a)  to  determine  whether  the  tendency  of  a 
typical  organic  oxidation  from  one  non-electrolyte  to  another,   or  its 
reversal,  to  take  place  gives  rise  to  a  constant,  measurable  and  reproduci- 
ble potential  difference,  when  the  concentrations  of  all  of  the  substances 
entering  into  the  electrochemical  equation,    by  which  the  process  may 
be  theoretically  represented,  are  fixed;  and  (b)  to  determine  whether  this 
potential  varies  with  the  concentrations  of  these  substances  in  the  same 
way  that  inorganic  potentials  have  been  found  to  vary,  viz.,  approxi- 
mately in  accordance  with  the  law  of  Mass  Action  as  expressed  in  the 
van't  Hoff  isotherm. 

(2)  A  typical  case  was  chosen  which  was  practically  adapted  to  such 
an  investigation,  namely,  the  oxidation  of  hydroquinone  to  quinone,  or 
its  reverse.     It  was  found  that  whenever  the  concentrations  could  be 
fixed  or  determined  to  a  reasonable  degree  of  certainty,  the  deviation 
from  the  van't  Hoff  equation  was  less  than  the  average  of  the  deviations 
obtained  in  parallel  work  on  inorganic  electrolytes. 

(3)  It  has  been  shown  experimentally  in  the  specific  instance  investi- 
gated that  the  electrochemical  theory  of  oxidation  and  reduction,  in  its 
full  quantitative  significance,  is  applicable  to  organic  non-electrolytes  in 
the  same  way  and  to  the  same  extent  as  to  inorganic  electrolytes;  that 
there  is  no  fundamental  difference  between  the  two  classes  of  substances, 
in  this  respect;  that  nothing  but  practical  difficulties,  peculiar  to  a  par- 
ticular case  in  our  opinion  stand  in  the  way  of  extending  these  relation- 
ships to  other  organic  systems. 

Incidentally,  it  has  been  shown  that  the  solubility  product  law,  which 
has  been  developed  in  connection  with  the  ionic  dissociation  of  difficultly 
soluble  electrolytes  and  has  already  been  applied  to  quinhydrone  by 
Luther  and  Leubner,  although  apparently  applying  with  a  fair  degree 
of  accuracy  to  the  dissociation  of  quinhydrone  in  the  presence  of  small 
excesses  of  hydroquinone,  breaks  down,  evidently,  near  the  point  at 
which  the  solution  is  saturated  with  respect  to  hydroquinone,  and  is 
actually  reversed  at  the  saturation  point;  and  that  this  reversal  is  prob- 
ably due,  not  to  a  change  in  the  true  solubility  of  the  quinhydrone,  but  to 
a  rapid  increase  in  the  dissociation  constant  in  the  neighborhood  of  the 
saturation  point  for  hydroquinone. 


PRELIMINARY  NOTE. 

A  certain  amount  of  freedom  has  been  exercised  in  this  work,  in  the  use  of  cer- 
tain terms. 

An  electrolyte  is  generally  defined  as  a  solution  which  conducts  electricity.  For 
lack  of  a  better  word,  however,  the  author  has  used  this  term  to  refer  to  the  dissolved 
substance  to  which  the  solution  owes  its  conductivity.  The  sense  in  which  the  term 
is  used  will  be  apparent  from  the  context.  No  terms  could  be  found  which  fitted  the 
case  exactly. 

Again,  the  author  has  occasionally  used  the  terms  potential,  oxidizing  potential 
and  reducing  potential  instead  of  the  usual  expression  potential  difference.  The  latter 
term,  in  this  connection,  has  a  purely  electrical  significance,  whereas,  when  the  author 
omits  the  word  "difference,"  he  is  referring  to  a  different  thing,  although  it  may  have 
the  same  numerical  value.  He  is  using  the  term  "potential"  in  a  chemical,  rather  than 
in  a  purely  electrical  sense.  In  either  case,  however,  it  refers  to  the  so-called  "inten- 
sity" or  "head"  factor  of  the  form  of  potential  energy  in  question. 

In  this  work  the  directly  observed  electromotive  force  measurements  are,  for  the 
sake  of  brevity,  omitted.  Only  the  so-called  single  or  pole  potential  differences  are 
given.  Neglecting  the  unknown  potential  difference  at  the  boundary  between  the  two 
solutions  making  up  the  cell,  the  electromotive  force  of  the  cell,  in  the  direction  from 
organic  electrode  through  external  circuit  to  calomel  electrode,  may  be  regarded  as 
the  difference  between  the  organic  pole  potential  difference  and  the  calomel  pole  poten- 
tial difference,  both  taken  in  the  direction  from  pole  to  solution.  The  organic  pole 
potential  difference  is  therefore  derived  by  algebraically  adding  the  calomel  pole  poten- 
tial difference  value,  taken  as  standard,  to  the  observed  electromotive  force — all  values 
being  taken  in  the  directions  above  mentioned. 

The  pole  polential  difference,  for  a  given  system  measures  the  so-called  oxidizing 
"power"  of  that  system  since  it  measures  its  tendency  to  give  up  positive  electricity. 
These  oxidizing  "powers,"  expressed  in  this  electrical  way,  are  sometimes  called  oxi- 
dizing potentials  or  simply  potentials,  because  of  their  analogy  to  electrical  and  mechanical 
po  entials,  and  because  they  can  be  balanced  by  electrical  potential  d  fferences,  by 
which  means  they  are  measured,  and  therefore  expressed  in  he  s:me  units.  Oxi- 
dizing potential  differs  from  e!ectrical  potential  difference,  however,  in  that  it  is  an  inde- 
pendent specific  property  of  the  system  in  question  for  the  given  concentrations  and 
temperature,  involving  nothing  external  to  the  system,  instead  of  a  relation  between 
it  and  its  environment,  imparted  to  it  by  temporary  external  conditions.  It  is  there- 
fore spoken  of  as  a  potential  rather  than  as  a  potential  difference.  The  relati  nship 
between  oxidizing  "power"  or  potential  and  the  electrical  potential  difference  by  which 
it  is  measured  is  analogous  to  the  relationship  between  the  pressure  of  a  steam  boiler, 
which  is  an  independent  property  of  the  system,  wet  steam,  for  the  conditions  existing 
within  the  boiler,  and  the  mercury  level  difference  of  a  manometer  by  which  it  might 
be  measured. 

In  the  present  work  we  are  concerned  with  the  chemical  significance  of  the  values  in 
question,  rather  than  with  their  electrical  significance.  When  the  term  potential  is 
used  alone,  in  this  work,  instead  of  potential  difference,  it  is  used  in  the  chemical  sense — 
in  the  sense  of  oxidizing  potential.  The  reducing  potential  of  a  system,  i.  e.,  its  ten- 
dency to  give  up  negative  electricity,  is  simply  its  oxidizing  potential  with  the  sign 
changed. 


INTRODUCTION. 

It  is  customary  at  the  present  time,  as  is  evident  from  the  contents 
of  our  text-books  on  organic  chemistry,  to  treat  the  subject  of  oxidation- 
reduction  of  organic  compounds  in  a  manner  similar  to  that  followed  in 
the  field  of  inorganic  chemistry  previous  to  the  electro-chemical  method. 
In  other  words,  more  attention  is  directed  towards  the  selection  of  the 
most  suitable  oxidizing  or  reducing  agent  for  the  particular  reaction  at 
hand,  than  to  a  consideration  of  the  dynamics  or  energy  relationships 
involved.  The  oxidizing  agent  and  the  temperature,  acidity,  alkalinity 
and  concentrations  of  the  solution,  etc.,  which  furnish  the  maximum 
yield  of  the  oxidation  product,  have  been  selected  from  trial  experiments 
rather  than  from  reasoning  based  upon  recognized  general  principles 
underlying  reactions  of  this  kind.  Furthermore,  some  claim1  that  this 
electrochemical  method  of  studying  the  energy  changes  in  reactions  like 
this  is  only  applicable  .to  electrolytes,  and  that  it  is  entirely  too  hypo- 
thetical to  view  oxidation  and  reduction,  in  the  case  of  organic  com- 
pounds, in  this  way.  It  is  quite  generally  accepted  that  the  electro- 
chemical way  for  accounting  for  oxidation  and  reduction,  in  the  case  of 
electrolytes,  is  far  superior  to  the  older  way,  based  upon  such  ideas  as  sub- 
stitution of  oxygen  for  hydrogen,  removal  of  hydrogen,  addition  of  oxy- 
gen, etc.  This  is  very  often  emphasized  in  the  lecture  room  by  the  familiar 
experiment  known  as  "oxidation  at  a  distance."2  It  therefore  follows 
that  if  it  is  not  permissible,  until  more  experimental  data  are  available, 
to  use  the  electrochemical  way  for  explaining  oxidation  and  reduction  in 
the  case  of  non-electrolytes,  the  older  way  must  be  employed.  Thus  we 
will  have  two  theories  for  oxidation-reduction,  one  for  inorganic  and  one 
for  organic  chemistry. 

PREVIOUS   WORK   IN   ORGANIC   CHEMISTRY. 

In  the  field  of  organic  chemistry,  the  very  little  work  that  has  been  done 
at  all,  in  the  investigation  of  Oxidizing  Potentials,  has  been  only  of  the 
most  empirical  character.  Bancroft,  in  i892,3  included,  among  the  large 
number  of  more  or  less  common  oxidizing  and  reducing  mixtures  with 
which  he  worked,  two  organic  reducing  agents,  namely,  alkaline  solu- 
tions of  hydroquinone  and  of  pyrogallol,  of  indefinite  composition.  The 
proportions  used  in  making  up  the  organic  solutions  were  not  given,  but 
had  they  been,  they  would  not  have  thrown  much  light  upon  the  composi- 
tion of  the  resulting  mixture  since  hydroquinone  and  pyrogallol  not  only 

1  Bray  and  Branch,  /.  Am.  Chem.  Soc.,  35,  1440  (1913);  Lewis,  Ibid.,  35, 1448  (1913) . 

2  See  "Qualitative  Chemical  Analysis,"  Stieglitz,  p.  253. 

3  Z.  physik.  Ch.,  10,  387. 


10 

an  electronic  transfer  which  can  be  used  to  develop  an  e.  m.  f . ;  as  the  pre- 
liminary study  is  certainly  in  harmony  with  this  idea  we  may  interpret 
the  reaction  as  follows  for  the  present: 

CH2  =  CH2  +  C12  — ^  +CH2—  +CH2  +  2C1-   ->  CH2C1— CH2C1." 
In  the  second,  a  similar  experiment  was  carried  out  with  the  cell 

HCHO  |  AT/5  NaOH  |  O2. 

Before  adding  the  formaldehyde,  oxygen  was  bubbled  against  one  of  the 
platinized  electrodes,  in  the  U-tube  of  sodium  hydroxide  solution,  and  the 
electromotive  force  measured.  Every  time  a  drop  of  formalin  solution 
was  added  to  the  other  arm,  the  electromotive  force  would  show  an  abrupt 
increase  to  a  comparatively  constant  value.  Similar  experiments  were 
carried  out  with  phenolphthalin  (reduction  product  of  phenolphthalein) 
and  oxygen  and  with  phenolphthalein  and  hydrogen.  In  the  first  case  no 
perceptible  pink  color  appeared  on  connecting  the  poles  until  a  current 
was  forced  through  the  cell.  In  the  second  case,  however,  the  solution 
became  lighter  without  the  use  of  an  external  current. 

This  concludes  the  list  of  all  the  work  that  could  be  found  on  organic 
potentials  up  to  the  present.  If  we  compare  this  with  what  has  been  done 
in  inorganic  chemistry,  the  contrast  is  apparent.  In  inorganic  chem- 
istry, investigation  has  invariably  shown  that  potentials  varied  with  the 
concentrations  of  the  substances  involved  in  the  electrochemical  reaction, 
to  which  the  oxidation  or  reduction  was  theoretically  attributable.  And 
when  the  absence  of  interfering  factors  seemed  to  permit,  which  was  true 
in  nearly  every  case  investigated,  this  variation  was  found  to  be  in  as  good 
accordance  as  could  be  expected  with  the  van't  Hoff  Isotherm,  which  is 
the  thermodynamical  verification  of  the  Law  of  Mass  Action.1 

Now,  in  every  one  of  these  cases,  either  the  oxidizing  or  reducing  agent, 
as  the  case  may  be,  or  its  product,  was  an  electrolyte.  The  van't  Hoff 
equation  has  never  been  applied  to  a  case  where  both  were  non-electro- 
lytes. In  none  of  the  organic  cases  just  reviewed,  has  any  attempt  been 
made  to  apply  this  or  any  other  generalization,  or  to  find  any  new  rela- 
tionship; to  establish  any  connection  between  the  potentials  measured 
and  any  particular  reaction  or  principle.  In  all  of  the  organic  work,  no 
account  has  been  taken  of  the  products  of  the  reaction.  There  has  been 
no  attempt  at  coordination  of  any  kind.  They  are  simply  isolated  and 
empirical  measurements. 

1  Nernst,  "Theoretical  Chemistry"  (English  Ed.),  p.  690;  Ibid.,  p.  785;  Peters,  Z. 
physik.  Ch.,  26,  193  (1898);  Friedenhagen,  Z.  anorg.  Ch.,  29,  396  (1902);  Tower,  Z. 
physik.  Ch.,  18,  17  (1.895);  Luther  and  Michie,  Z.  Electroch.,  14,  126  (1908);  Luther  and 
Sammet,  Ibid.,  n,  293  (1905).  Since  this  equation  is  based  on  the  perfect  gas  laws,  or, 
in  the  case  of  solutions,  on  the  corresponding  relationships  for  osmotic  pressure,  it  is 
only  applicable  to  the  degree  in  which  these  laws  hold,  in  the  case  in  question. 


T  I 


According  to  the  van't  Hoff  equation  or  the  principles  of  chemical 
equilibrium,  a  reducing  or  oxidizing  agent,  in  solution  absolutely  free  from 
any  of  its  product,  should  have  a  potential  of  infinity  (minus  or  plus). 
Therefore,  since  so-called  "pure"  solutions  of  oxidizing  or  reducing  agents 
do  not  show  infinite  potentials,  the  presumption  is  that  they  contain 
quantities  of  products  sufficient  to  account  for  the  potentials  observed, 
though  too  small  to  be  detected  by  analytical  means.  Actual  calcula- 
tions from  existing  experimental  data  will  show  that  the  concentrations 
of  these  substances  can  be  that  small  without  the  theoretical  potentials 
being  any  larger  than  those  actually  found  for  pure  solutions,  which  are 
usually  of  the  order  of  about  i  volt,  and  never  more  than  two.  This  is 
reasonable  enough,  for,  if  the  solution  were  so  pure  as  to  give  a  potential 
greater  than  this,  its  oxidizing  or  reducing  power  would  be  so  great  that, 
we  would  expect  it  to  react  with  the  water,  or  anything  else  available 
producing  product  until  the  excessive  tension  was  relieved.  Measure- 
ments on  so-called  pure  agents  are  therefore  to  be  taken,  as  quantitative 
data,  with  more  or  less  reserve,  as  we  are  not  justified  in  assuming  that 
these  indeterminately  small  quantities  of  products  present  may  not  vary 
greatly,  relatively  speaking,  in  different  specimens,  even  though  a  par- 
ticular series  of  measurements  has  given  fairly  reproducible  results.  One 
more  point  may  be  mentioned.  Mere  potential  measurements,  without 
a  correlation  of  some  sort,  prove  nothing  even  qualitatively.  Any  two 
solutions  containing  electrolytes  'will  give  a  potential  difference,  due  to 
the  relative  discharging  tendencies  of  their  ions,  even  when  no  oxidizing 
or  reducing  agent  is  added,  unless,  by  a  rare  coincidence  they  happen  to 
have  practically  the  same  oxidizing  potential. 

PURPOSE    OF   PRESENT   WORK. 

From  what  has  just  been  said  it  'is  evident  that  we  cannot  consider 
a  potential  measurement  made  upon  a  "pure"  solution  of  a  reducing 
agent  as  a  characteristic  fixed  property  of  that  substance  (like,  for  in- 
stance, a  melting  point).  Such  a  value,  on  the  contrary,  would  be  an 
accidental  one.  If  it  is  only  desired  to  measure  the  reducing  power  of 
that  particular  solution,  for  empirical  purposes,  then  such  a  measurement 
gives  all  the  information  that  is  required.  But  it  is  useless  as  a  basis 
for  general  calculations.  For  it  has  been  found,  whenever  investigated, 
that  the  potential  is  as  much  a  property  of  the  oxidation  product  of  the 
reducing  agent  as  of  the  reducing  agent  itself,  and  is  a  function  of  the  con- 
centrations of  both.  In  other  words,  it  is  a  property  of  the  system  rather 
than  of  the  substance.  There  is,  of  course,  such  a  thing  as  the  intrinsic 
tendency  of  the  oxidizing  or  reducing  agent  to  pass  to  another  stage  of 
oxidation,  by  virtue  of  which  it  owes  its  oxidizing  or  reducing  properties. 
This  is  a  characteristic  property  of  the  substance,  being  independent  of 


12 


concentrations  and  represented  by  the  term  — RT/wF  logeK0   in  the 
van't  Hoff  equation,  in  the  form: 

RT         C,W'CV"..  RT 


*  =   ~^T  log,  n  ,ni,n  ,n*r  ~^T  iogeKo  or 

\^i      \^z       •  •  •        Mr 

RT         CimC2W2 

7T    =    —    loge  c   m/c  /Bj  '     -    +    7T0.  (l) 

It  determines  the  equilibrium  constant.  But  we  cannot  measure  this 
directly,  by  any  means  known  to  us  at  the  present  time.  We  can  only 
determine  the  potential  of  the  system,  under  known  concentration  con- 
ditions, and  for  this  calculate  by  means  of  the  above  equation  the  intrinsic 
oxidizing  of  reducing  tendency,  which  is  sometimes  written  TTO  and  some- 
times A.1  We  must  first,  therefore,  know  all  the  concentrations  involved 
and  test  our  relationship. 

These  facts  have  been  recognized  in  inorganic  chemistry  for  years 
(though  confused  by  the  tendency  to  attribute  all  oxidation  to  oxygen 
and  all  reduction  to  hydrogen).  In  organic  chemistry,  however,  as  has 
been  shown,  no  attempt  has  ever  been  made,  so  far  as  we  can  discover, 
to  put  organic  oxidation  and  reduction  on  a  similar  basis.  The  idea 
seems  to  have  prevailed  that  the  electrochemical  theory  of  oxidation 
and  reduction  was  not  applicable  to  non-electrolytic  substances,  such 
substances  being  oxidized  or  reduced  through  the  medium  of  "nascent" 
oxygen  or  hydrogen  liberated  by  the  oxidizing  or  reducing  agent,  or  by 
the  removal  of  hydrogen  or  oxygen,  from  the  non-electrolytic  substance 
by  the  agent.  The  purpose  of  the  present  investigation  is  to  apply  the 
electrochemical  theory  to  a  typical  organic  oxidation-reduction  in  a  man- 
ner parallel  to  that  in  which  it  has  been  applied  in  inorganic  chemistry, 
and  to  see  whether  or  not  the  theory  fits  the  observed  facts  as  satisfac- 
torily as  it  has  been  found  to  in  inorganic  chemistry. 

The  conception  that  an  electric  charge  can  be  given  up  by  an  elec- 
trically neutral  body  and  still  leave  an  electrically  neutral  body  is,  of 
course  easily  explained  by  the  giving  up  of  an  oppositely  charged  ion  at 
the  same  time,  as  illustrated  by  the  following  two  equations: 

CH3CH3OH  ^  CH3CHO  +  20  +  2H+  (2) 

C6H4(OH)2  -  C6H402  +  20  +  2H+  (3) 

SELECTION   OF  A   SYSTEM. 

The  first  problem  confronted  was  the  selection  of  a  suitable  system, 
one  that  would  adapt  itself  to  such  an  investigation.  The  required  quali- 
fications were  the  following: 

(i)  It  must  be  a  typical  reversible  organic  oxidation. 
1  This  should  not  be  confused  with  the  use  of  the  symbol  A  to  represent  free  energy. 


13 

(2)  The  reaction  must  be  well  defined,  that  is,  the  equation  must  be 
definitely  known,  and  it  must  be  as  free  as  possible  from  intermediate 
stages  and  tendency  to  go  to  further  stages  of  oxidation. 

(3)  It  must  be  possible  to  determine  the  concentrations  of  the  sub- 
stances  involved,   with  sufficient  accuracy,  by  analytical  means,  or  else 
to  fix  or  control  them  in  some  way. 

(4)  The  substances  involved  must  all  be  sufficiently  stable  so  that  the 
system  will  not  become  lost  in  a  maze  of  unknown  side  reactions  before 
a  definite  potential  can  be  obtained.     And  there  must  be  no  reactions 
entering  in  to  throw  the  potential  off  and  prevent  the  true  reading  from 
being  obtained. 

(5)  The  substances  must  all  be  sufficiently  non-volatile  so  as  not  to  be 
removed  appreciably  by  the  bubbling  of  an  inert  gas  through  the  solution 
to  remove  the  air. 

(6)  They  must  be  appreciably  soluble  in  water. 

(7)  They  should  not  be  acids  or  bases  or  react  in  any  way  with  the  acid 
or  base  used  for  fixing  the  hydrogen  ion  concentration,  unless  the  concen- 
trations can  still  be  controlled,  in  spite  of  this  fact. 

It  is  evident  that  a  system  fulfilling  these  specifications  is  not  easy  to 
find  in  organic  chemistry. 

The  system  hydroquinone/quinone  was  selected  as  more  nearly  meet- 
ing these  requirements  than  any  other  system  that  was  thought  of.  Its 
advantages  lay  in  the  following  facts:  The  solubilities  were  such  that 
the  concentrations  could  be  fixed  by  using  saturated  solutions,  which  would 
still  be  fairly  dilute.  It  is  a  well  defined,  clean  cut,  perfectly  reversible 
reaction.  Neither  substance  is  very  volatile.  Quinone  is  about  as  vola- 
tile as  water,  but  the  working  conditions  were  not  such  that  this  gave 
any  trouble.  Neither  substance  is  a  pronounced  acid  or  base.  The  reac- 
tion is  a  rapid  one.  The  substances  are  not  very  unstable,  except  in  alka- 
line solution.  Quinone  can  be  determined  accurately  by  analytical 
means,  even  in  the  presence  of  hydroquinone.  The  system  possesses, 
however,  certain  disadvantages.  Hydroquinone,  when  added  to  an  alka- 
line solution,  is  a  strong  enough  acid  to  nearly  neutralize  the  alkali,  so 
that  a  special  investigation  would  have  to  be  undertaken  to  find  a  means 
of  accurately  determining  the  hydrogen  ion  and  other  concentrations. 
Furthermore,  hydroquinone  and  quinone  are  both  very  unstable  in  alka- 
line solution,  somewhat  so  even  in  pure  water  and  more  so  in  neutral 
potassium  chloride  solution,  the  solution  becoming  dark  brown  owing  to 
the  formation  of  a  tarry  material.  Finally,  the  system  is  complicated 
by  the  fact  that  hydroquinone  and  quinone  combine  directly,  in  solution 
and  elsewhere,  to  form  the  beautiful  bronze-green  crystalline  addition 
product,  quinhydrone. 


QUINHYDRONE. 

Quinhydrone  is  formed  immediately  whenever  quinone  is  added  to  a 
solution  of  hydroquinone,  and  vice  versa.  In  solution  it  exists  in  mobile 
equilibrium  with  its  components.  In  aqueous  solution,  it  is  highly  dis- 
sociated (over  90%),  but  its  solubility  is  so  slight  that  it  is  impossible 
for  a  considerable  concentration  of  one  of  its  constituent  substances  to 
exist  in  the  presence  of  a  considerable  concentration  of  the  other.  There- 
fore, a  solution  cannot  be  simultaneously  saturated  with  hydroquinone 
and  quinone,  the  molar  solubilities  of  these  substances  being,  respec- 
tively, about  640  and  125  times  that  of  quinhy drone.  In  a  nearly  satura- 
ted solution  of  either,  the  addition  of  only  a  small  quantity  of  the  other  is 
required  to  produce  a  precipitate  of  quinhydrone. 

INFLUENCE   OF   HYDROGEN   ION   CONCENTRATION. 

In  Equations  (2)  and  (3)  given  above,  hydrogen  ion  appears  as  a  prod- 
uct of  the  oxidation  of  alcohol  and  of  hydroquinone,  and  would  appear, 
similarly,  in  the  cases  of  most  organic  oxidations.  Theoretically,  then, 
acidity  should  favor  the  reaction  in  the  oxidizing  direction  (backwards, 
as  written)  and  oppose  it  in  the  reducing  direction.  The  reducing  poten- 
tial of  the  system  should  therefore  decrease,  that  is,  the  oxidizing  poten- 
tial should  increase,  with  increasing  hydrogen-ion  concentration. 

Acidity  and  alkalinity  have  long  been  recognized  as  important  fac- 
tors in  determining  the  course  of  oxidation  and  reduction,  especially  in 
organic  chemistry.  It  would  therefore  seem  of  importance  to  test  the 
above  theoretical  influence  experimentally.  This  never  seems  to  have 
been  done,  even  qualitatively,  in  organic  chemistry,  though  the  quantita- 
tive theory  has  been  checked  in  a  number  of  cases  in  inorganic  chemis- 
try. 

But  because  it  is  found  that  the  oxidizing  potential  of  the  system  is  favored  by 
acidity,  and  the  reducing  potential  by  alkalinity,  it  does  not  follow  that  alkalinity  should 
always  be  found  to  favor  the  action  of  hydroquinone  as  a  reducing  agent  or  acidity  to 
favor  in  every  case  the  reduction  of  quinone.  The  influence  of  hydrogen  ion  upon  the 
other  system,  entering  into  the  reaction,  must  be  taken  into  account.  For  instance, 
the  oxidation  of  hydroquinone  by  a  ferric  salt  should  be  hindered  and  the  reduction  of 
quinone  by  an  iodide,  favored  by  acidity,  for  we  have: 

C6H4(OH)2  ^±  C6H402  +  20  +  2H+     . 

20  +  2Fe  +  +  +^  2Fe^ 


C6H4(OH)2  +  2Fe+++  ^  C6H402  +  2Fe  +  + 
and 

C6H402  +  2H+  ^  C6H4(OH)2  +  20 
20  +  2l~  —  I2 


CGH402  +  2H+  +  2l~  —  C6H4(OH) 


(4) 


(5) 


15 

But  the  oxidation  of  hydroquinone  to  quinone  by  permanganate  should  be  favored  by 
acidity,  because 

5(C6H4(OH)2  ^  C6H402  +  2H+  +  20) 

__  2(MnQ4-  +  8H  +  —  Mn++  +  50  +  4H2O) 
5C6H4(OH)2  -f  2Mn04-  +  6H+  ^  5C6H4O2  +  2Mn++  +  8H2O 

In  the  hydroquinone  reaction,  it  will  be  seen,  we  have  involved  only  one  hydro- 
gen ion  per  charge  transferred,  whereas  in  the  permanganate  reaction  the  ratio  is  8  to 
5,  so  that  the  influence  in  the  latter  system  outweighs  that  in  the  former. 

It  makes  a  difference  also  what  concentrations  are  being  controlled.  For  instance 
the  oxidation  of  chromium  from  the  trivalent  to  the  hexavalent  condition  by  sodium 
peroxide  may  be  written  so  as  to  appear  to  be  either  hindered  or  favo  ed  by  alkalinity. 

2(Cr(OH)3  +  50H-  ^±  CrO4=  +  4H2O  +30) 
3(2H20  +  O,"  ^±  4QH-  +  2®) 


2Cr(OH)3  +  3O2  =  ^  2CrO4=  +  2H2O  +  2OH~ 
or 

2(Cr  +  +  +  +  80H-  ^±  Cr04=  +  4H2O  +  3©) 

3(2H20  +  02=  ^  4OH-  +  20) 


-  4OH~  +  3O2=  ^  2CrO4~  +  2H2O 

This  simply  means  that,  if  we  are  starting  with  chromic  hydroxide,  excess  alkali  would 
tend  to  oppose  the  oxidation,  alkali  being  one  of  the  products ;  whereas,  if  we  are  starting 
with  a  chromic  salt,  additional  alkali  would  be  used  up  in  the  formation  of  the  hy- 
droxide. (This  may  not  be  true,  experimentally.  Other  factors,  which  we  have  not 
considered,  which  are  unknown  perhaps,  may  enter  in.  For  instance,  acidity  or  de- 
creasing alkalinity  would  hinder  the  reaction  by  removing  peroxide  ion.  But  the  above 
would  be  the  interpretation  of  the  equations  as  they  stand  )  So  we  must  consider  the 
actual  conditions  in  writing  our  oxidation  equations. 

In  order  to  determine  the  influence  of  hydrogen-ion  concentration  on 
the  potential,  it  was  hoped  that  a  system  could  be  found  in  which  the 
concentrations  of  the  other  factors  could  be  fixed  by  the  use  of  saturated 
solutions  of  the  substances  involved,  while  the  acidity  varied.  In  the  one 
selected,  however,  the  formation  of  quinhydrone  made  this  impossible, 
as  already  pointed  out.  But,  fortunately,  the  concentrations  could  be 
fixed,  as  well,  though  indirectly,  by  saturating  the  solution  with  quin- 
hydrone and  one  of  the  components. 

PLAN   OF   THE   WORK. 

In  view  of  the  considerations  outlined  above,  the  investigation  was 
carried  out  according  to  the  following  plan:  Potential  measurements 
were  made  on  a  series  of  solutions  saturated  simultaneously  with  hydro- 
quinone and  quinhydrone,  in  which  the  hydrogen-ion  concentration  was 
varied  and  the  conductivity  furnished  by  the  use  of  hydrochloric  acid 
and  sodium  hydroxide,  in  various  concentrations;  for  the  neutral  solu- 
tion, potassium  chloride  was  added  as  electrolyte.  A  parallel  series  was 
started,  with  quinone  and  quinhydrone,  but  because  of  the  instability 
of  the  former,  it  was  not  carried  any  further  than  normal  and  tenth- 


i6 

normal  acid.  To  determine  the  effect  of  varying  other  concentrations, 
a  series  was  run  in  tenth-normal  acid  saturated  with  quinhydrone  with 
varying  concentrations  of  hydroquinone  added.  In  order  to  know  the 
concentrations  it  was  necessary  to  determine  the  solubilities  of  the  sub- 
stances and  the  dissociation  constant  of  quinhydrone.  The  results  were 
compared  with  the  van't  Hoff  equation,  as  a  convenient  standard  of 
reference,  in  the  form 

RT ,       (quinone)  (H+)2       RT  t 
=  2-F  10g<  (hydroquinone)  '  ~  2T  ^  K° '  (9) 

In  neutral  and  alkaline  solutions,  because  of  the  instability  of  hydro- 
quinone and  quinone  and  other  difficulties,  data  of  definite  quantitative 
significance  were  not  obtained.  The  results  were,  however,  of  qualitative 
interest.  In  acid  solutions,  the  results  paralleled  those  in  inorganic  chem- 
istry. 


PART  I 

I.    POTENTIAL  MEASUREMENTS   IN   ACID   SOLUTIONS. 

Mixtures  were  made  up  as  described  below,  and  their  potentials  mea- 
sured in  a  half -element  vessel,  against  a  saturated  potassium  chloride 
calomel  electrode,  by  means  of  an  e.  m.  f.  combination  of  the  type 

Hg  |  HgCl  sat.  KC1  |  sat.  KC1  |  solution  A  |  Pt, 

employing  a  sensitive  potentiometer  and  galvanometer.  Saturated 
potassium  chloride  solution  was  used  as  the  connecting  medium.  The 
cells  were  kept  immersed  in  a  constant  temperature  bath  at  25°  C., 
the  temperature  remaining  constant  to  a  hundredth  of  a  degree. 

Nitrogen  was  bubbled  through  the  mixture  in  the  cell,  for  the  first  few 
hours,  to  insure  complete  removal  of  the  air  and  to  provide  agitation  at 
the  start.  The  gas  inlet  and  outlet  tubes  were  then  closed  to  prevent 
access  of  air.  The  potential  of  each  cell  was  measured  at  least  once  a 
day  over  a  period  of  from  one  to  three  weeks. 

The  following  mixtures  were  investigated: 

A.  Normal  hydrochloric  acid,  saturated  with  hydroquinone  and  quin- 
hydrone. 

B.  Tenth-normal  hydrochloric  acid,  saturated  with  hydroquinone  and 
quinhy  drone. 

C.  Hundredth-normal  hydrochloric  acid,  saturated  with  hydroquinone 
and  quinhydrone. 

D.  Normal  hydrochloric  acid,  saturated  with  quinone  and  quinhydrone. 
B.  Tenth-normal  hydrochloric  acid,  saturated  with  quinone  and  quin- 
hydrone. 

F.  A/"/ 10  HC1,  o.i  molar  hydroquinone,  saturated  quinhydrone. 

G.  N/io  HC1,  0.05  molar  hydroquinone,  saturated  quinhydrone. 
H.  A/YIO  HC1,  0.02  molar  hydroquinone,  saturated  quinhydrone. 
I.  N/io  HC1,  o.oi  molar  hydroquinone,  saturated  quinhydrone. 
J.  N/io  HC1,  saturated  with  quinhydrone  alone. 

Two  or  more  cells  of  each  mixture  were  made  up  and  examined,  usually 
at  different  times,  and  sometimes  with  different  lots  of  materials,  in 
order  to  determine  the  reproducibility  of  the  potentials  found. 

The  calomel  electrode  was  assigned  the  value  0.5265  volts,  which 
is  the  value  provisionally  adopted  by  Dr.  H.  A.  Fales1  for  25°  C. 
Adding  this  to  the  electromotive  force  of  the  cell,  when  the  mix- 
ture is  the  positive  electrode,  gives  the  single  potential  difference2 
or  the  oxidizing  potential  of  the  mixture  in  question,  plus  or  minus  any 
contact  potential  differences  at  the  boundaries  of  the  solutions.  The 

1  Private  communication. 

2  See  Le  Blanc,  "Text-Book  of  Electrochemistry,"  Eng.  Trans,  of  4th  Ger.  Ed.,  p. 
234,  et  seq.;  and  the  articles  already  cited. 


i8 


usual  custom  of  writing  potentials  with  a  minus  sign  before  them  when 
the  tendency  is  for  the  solution  to  give  up  positive  electricity  to  a  zero 
electrode,  i.  e.,  to  oxidize  it,  has  been  departed  from  because  it  seems  more 
logical  and  less  confusing  to  regard  such  potentials  as  positive.  That  is, 
such  a  mixture  would  be  said  to  have  a  positive  oxidizing  potential  or  a 
negative  reducing  potential,  and  we  are  expressing  the  values  as  oxidizing 
potentials.  The  question  is,  of  course,  an  arbitrary  one. 
The  results  are  given  in  Table  I. 

TABLE  I. 
MIXTURE  A. 


Cell  No.  1. 

Cell  No.  2. 

Cell  No.  3. 

Hours  from 

start.                Volts,  i 

Hours. 

Volts. 

Hours. 

Volts. 

o               0.8884 

iVi 

0.8889 

I23/4 

0.8834 

24l/z           0.8833 

23!/4 

0.8866 

293/4 

0.8832 

49J/4              0.8810 

5  iVi 

o  .  8847 

3I3A 

0.8830 

78                    0.8812 

77 

0.8871 

583A 

0.8830 

97  */4           0.8807 

953A 

0.8867 

7574 

0.8832 

I2lV2                0.8827 

ii93A 

0.8851 

9472 

0.8831 

143  V4      .          0.8805 

147 

0.8841 

II974 

0.8828 

16272                 0.8801 

16774 

0.8817 

H472 

0.8823 

Cell  No.  3   (Continued).                                                         I731/2 

o  .  8804 

Hours.                    Volts. 

I923/4 

0.8797 

3693/4                   0. 

8776 

21574 

0.8819 

384                    o. 

8769 

24l74 

0.88II 

40774                o. 

8767 

2653/4 

o  .  8805 

43274                       0. 

8765 

28972 

0.8786 

547                     o. 

8753 

31274 

0.8790 

479V2                o. 

8746 

33972 

0.8782 

MIXTURE  B. 

Cell  No.  1.                          Cell 

No.  2. 

Cell 

No.  3.                           Cell 

No.  4. 

Hours.           Volts.              Hours. 

Volts. 

Hours. 

Volts.             Hours. 

Volts. 

l/4                 0.8213              I8V4 

o  .  8309 

2'/2 

o  .  8303             o 

o.  8091 

29x/2                0.8234             24J/4 

o  .  8303 

21 

0.8306             872 

0.8293 

5lV2                0.82892           42J/4 

0.8319 

548/4 

0.8315                 22V2 

0.8299 

74J/2           0.82992       72 

0.82762 

693A 

0.8309           4772 

0.8296 

97                o.  83282        89J/2 

0.8282 

9i  3A 

0.8317          70 

0.8295 

I2IJ/4                  0.83I52 

i  M'A 

0.8311            93 

0.8289 

143 

O.83II              I2l72 

0.8285 

16774 

0.8303             I423/4 

0.8290 

Cell  No.  4  (Continued). 

i873/4 

O.83OO             16674 

0.8298 

Hours.             Volts. 

2IIl/2 

0.8294          19° 

0.8298 

338                  0.8297 

2353A 

0.     289             21474 

0.8294 

362  1/4              O.8295 

26972 

0.8288             24572 

o  .  8303 

386V2          0.8304 

2973A 

O.8266             264 

0.8300 

4I3V4          0.8292 

31372 

O.8274             29274 

0.8306 

432  V2               0.8278 

331 

0.8250             3097'2 

o  .  8303 

1  All  potential  values  given  in  this  table  are  the  single  potential  differences  de- 
rived from  the  observed  electromotive  force  of  the  combination  by  adding  the  value 
(0.5265)  of  the  calomel  electrode. 

2  Average  of  several  readings  taken  that  day. 


TA: 

BUS  I  (Continued}. 

MIXTURE  C. 

Cell  No. 

1: 

Cell  No. 

2. 

Hours. 

Volts. 

Hours. 

Volts. 

l/4 

0-7321 

Va 

0.7753 

2O 

0.7626 

20V2 

0.7756 

43  3  A 

0.7736 

50 

0.7734 

73'A 

0.7718 

66»/4 

0.7709 

90 

0.7714 

9I*/2 

0.7673 

if4Va 

0.7691 

I40V2 

0.7697 

i633A 

0.7718 

i633A 

0.7700 

187 

0.7725 

i87V2 

o  .  7696 

2IOl/2 

0.7725 

2i7Va 

o  .  7687 

2403A 

0.7708 

234V4 

0.7685 

257V2 

0.7712 

262  3/4 

o  .  7686 

286 

0.7738 

aftvVi 

o  .  7670 

305  V4 

0.7741 

307 

o  .  7642 

33i3A 

o  .  7662 

361 

o  .  7646 

380 

0.7634 

402  3  A 

0.7638 

MIXTURE  D. 

Cell  No. 

4. 

Cell  No. 

5. 

Hours. 

Volts. 

Hours. 

Volts. 

O 

I  •  0353 

0 

I  .OIO2 

3 

I  .0190 

l/2 

I.  0137 

22l/a 

0.9928 

2O 

o  .  9944 

49 

0.9845 

23J/2 

0.9924 

763A 

0.9850 

42  3  A 

0.9829 

94V4 

0.9847 

71 

0.9836 

I20V2 

o  .  9840 

93'A 

0.9841 

146 

0.9838 

116 

o  .  9868 

1  68 

0.9826 

I4lV4 

o  .  9860 

195  Va 

0.9835 

I62V4 

o  .  9842 

2I4l/« 

0.9831 

I86V2 

o  .  9884 

2383/4 

0.9835 

217 

o  .  9869 

26l 

0.9786 

239l/4 

o  .  9849 

2893/4 

0-9791 

2593A 

0.9835 

3I2X/4 

0.9781 

275 

0.9834 

334*/a 

0.9826 

3593A 

0.9825 

38o3/4 

0.9853 

405V4 

0.9862 

435V2 

0.9869 

4573A 

0.9762 

478V2 

o  .  9808 

20 


TABLE  I  (Continued). 
MIXTURE  E. 


Cell  No. 


Hours. 

Volts. 

0 

0.9517 

I53/4 

o  .  9502 

40»/t 

0.9469 

653A 

0-9434 

943A 

0.9390 

114 

0.9379 

I36V4 

0-9371 

l623/4 

0.9348 

187 

0.9330 

233V2 

0.9299 

26ol/2 

0.9285 

291 

0.9284 

306V4 

0.9268 

329V2 

0.9280 

3533A 

0.9279 

378*A 

0.9275 

406  'A 

0.9266 

448  l  A 

0.9307 

473*/4 

0.9315 

492 

o  .  9307 

Cell  No. 

i. 

Hours. 

Volts. 

43A 

0.8695 

iSVz 

o  .  8705 

38 

0.8701 

6S'/4 

0.8695 

85 

0.8689 

H33/4 

0.8685 

I42V4 

0.8676 

167 

0.8666 

191 

0.8663 

238V2 

0.8625 

Cell  No. 

1. 

Hours. 

Volts.  ' 

I8V2 

0.8860 

42  1  A 

0.8850  . 

75JA 

0.8837 

903/4 

0.8828 

I233A 

0.8807 

I471/* 

0.8790 

173 

0.8785 

243 

0.8754 

3I5V4 

0.8715 

MIXTURE  F. 


MIXTURE  G. 


Hours. 
O 
9 

22  1/4 

473A 
76'A 

96 

H8'/2 

I443A 

169 

215'A 

273 


335  3A 
360 1/«' 
388V, 

43Q3A 
455  V4 


Cell.  No.  2. 


Hours. 

23/4 

15 

343A 
643A 
Si1/* 

H03/4 

I363A 


.Cell  No.  2. 


235 


Volts. 

0.9524 
0.9505 
0.9496 
o . 9456 

0.9412 
0.9395 
0.9385 

o . 9360 

0.9340 
0.9307 

0.9290 

0.9293 

0.9287 
0.9301 
0.9306 

0.9297 
0.9285 
0.9318 

0.9327 
0.9319 

o  9330 


Volts. 

o . 8702 

0.8689 
0.8687 
0.8692 
0.8681 

0.8686 
0.8658 
0.8671 
0.8663 
0.8635 


Hours. 
I53A 
393/4 

72'A 

88 

I203/4 

W/! 
309 


Cell  No.  2. 


Volts. 

0.8860 

0.8854 

0.8850 

O . 8849 

0.8839 

0.8829 

0.8818 

0.8790 

0.8760 


21 


TABLE  I  (Concluded). 


Hours. 


Cell  No.  1. 


72 
I043/4 

154 
293 

Hours. 

33'A 
49 

8 11/* 
105  XA 

I4I  3/4 

2OI 

2693/4 


Cell  No.  1. 


Hours. 

43A 

I8V4 


89l/s 
109 
"8V* 


Cell  No.  1 . 


174'A 

202  »/2 
226V4 
251 

298V2 


Volts. 

o . 9030 

0.9000 
0.8992 
0.8982 
0.8966 
0.8951 
0.8933 
0.8893 
0.8839 


Volts. 
0.9049 

o . 9068 
0.9062 

0.9051 
0.9033 

0.9022 
0.8998 

0.8975 


Volts. 

0.9155 
0.9142 
0.9132 
0.9116 
0.9100 

o  9099 
0.9079 
o . 9069 

0.9047 
0.9024 
0.9012 

0.8995 

0.8952 


MIXTURE  H. 


MIXTURE  I. 


MIXTURE  J. 


Hours. 

v< 

21 

54 
693A 

102  Vi 
126 

i5iVi 

I733A 
i94Vi 


Cell  No.  2. 


Volts. 

0.9018 

0.9014 

0.9007 

0.8998 

0.8987 

0.8978 

0.8969 

0.8938 

0.8921 


Hours. 

Vi 

303/4 

46'A 


Cell  No.  2. 


102  3/4 

"8V4 

i98Vi 
267 


Hours. 

I3V4 

373A 
69 

8?Vt 
ii63/4 

I34V4 
i64Vi 
i883/4 

211 


Cell  No.  2. 


Volts. 
0.9071 
0.9062 
o . 9062 

0.9055 
0.9044 
0.9027 

o . 8994 

0.8971 

Volts. 

0.9136 
0.9130 
0.9119 
0.9108 

0.9094 

o . 9082 

0.9067 
0.9054 

o . 9042 

0.9007 


From  the  above  data,  it  will  be  seen  that  each  of  the  various  mixtures 
has  its  characteristic  potential,  which  is  fairly  reproducible,  the  potentials 
being  obviously,  therefore,  a  function  of  the  relative  concentrations  of 
the  substances  involved,  and  that  the  potentials  given  by  the  solutions 
saturated  with  hydroquinone  become  "constant"  almost  from  the  start, 
and  remain  so  for  a  long  time,  after  which  they  begin  to  fall  off  gradually. 
The  hydroquinone  solutions  which  were  not  saturated  with  it,  however, 
showed  potentials  which  fell  off  almost  steadily  throughout  the  period 
during  which  the  cell  was  kept,  at  the  rate,  usually,  of  about  a  milli- 


22 

volt  per  day,  showing  the  necessity  of  fixing  the  concentrations  by  satura- 
tion, in  order  to  obtain  constant  potentials.  The  potentials  given  by  the 
solutions  saturated  with  quinone  showed  the  peculiar  behavior  of  falling 
several  hundredths  of  a  volt,  and  then  remaining  stationary  for  a  long  time. 
In  the  tenth-normal  acid,  this  constant  value  was  reached  only  after  a 
much  longer  time  than  in  the  case  of  the  normal  acid,  but  the  fall  in  the 
case  of  the  latter  was  greater.  Moreover,  it  was  only  when  a  very  large 
excess  of  solid  quinone  was  present,  that  these  stationary  periods  were 
obtained  with  D  and  B  at  all. 

The  above  points  and  the  exact  degree  of  constancy  and  reproducibility 
are  brought  out  more  clearly  by  graphical  representation  in  the  charts,  in 
which  the  above  potentials  are  plotted  against  time.1 

It  is  convenient  to  consider  the  above  data  in  groups  or  series.  'Series  I, 
consisting  of  Mixtures  A,  B,  C  and  Series  II,  consisting  of  Mixtures  D 
and  K,  show  the  effect  of  varying  hydrogen-ion  concentration  or  acidity 
on  the  potentials,  the  concentrations  of  the  other  substances  involved 
being  kept  approximately  constant,  in  each  series.  Series  III,  consisting 
of  mixtures  B,  F,  G,  H,  I,  J  and  K,  shows  the  effect  of  varying  hydro- 
quinone  and  quinone  concentrations  with  approximately  constant  acidity. 

II.    SOLUBILITIES   AND   THE   DISSOCIATION   OF   QUINHYDRONE. 

The  purpose  in  using  saturated  solutions  was  to  fix  the  concentrations 
of  the  substances  in  question,  throughout  the  life  of  the  cell,  regardless 
of  side  reactions  taking  place.  If  we  wish  to  know  what  these  concentra- 
tions are,  however,  we  must  know  the  solubilities  of  the  substances. 

Solubility  of  Hydroquinone.  —  No  data  on  the  solubility  of  hydroquinone 
at  25°  C.  could  be  found  in  the  literature,  so  it  was  determined  by  evap- 
orating 5  cc.  of  the  saturated  solution  to  constant  weight,  in  vacuum,  at 
room  temperature,  in  a  small  weighed  flask.  The  results  were  as  follows  : 

II. 


Grams  per  100  cc.  solution.  Moles 

--  2  --  .  -----         per 
Solvent.  No.  1.         No.  2.          No.  3.          No.  4.          No.  5.      Average.       liter. 

Water  .............  7-094  7  .091  7.112       7.086                       7.10  0.645 

o.oi  N  HC1  ........  7.060  7.128  7.136       7.028       7.146       7.10  0.645 

o.iNHCl  ..........  6.978  6.944                                                       6-96  0.633 

NHC1  .............  5.436  5.442                                                       5-44  0.494 

It  will  be  seen  that  the  solubility  of  hydroquinone  in  water  is  decidedly 
decreased  by  the  presence  of  hydrochloric  acid. 

Solubility  of  Quinone.  —  The  solubility  of  quinone  in  water  at  25°  C.  was 
determined  by  Robt.  Luther  and  A.  Leubner,2  using  the  analytical  method 
of  Amand  Valeur,3  which  consists  in  titrating  the  iodine  liberated  by 

1  See  pp.  29  and  32. 

2  /.  /.  prakt.  Chem.,  n.  f.,  85,  314-321  (1912). 

3  Compt.  rend.,  129,  552  (1899). 


23 

quinone  from  a  hydrochloric  acid  solution  of  potassium  iodide,  with  thio- 
sulphate.  They  give  the  value  1.265  moles  per  liter  in  which  there  is 
evidently  an  error  in  the  placing  of  the  decimal  point,  since  this  figure 
is  just  about  ten  times  too  large. 

Using  the  same  method,  we  obtained  the  following: 

TABLE  III. 

. Moles  per  liter. > 

Solvent.  No.  1.  No.  2.  Grams  per  100  cc. 

Water 0.1266  0.1266  1.37 

JV/ioHCl 0.1275  0.1275  1-38 

NHC1 ' 0.1332  0.1332  1-44 

It  will  be  observed  that  the  solubility  of  quinone  in  water  is  decidedly 
increased  by  the  presence  of  hydrochloric  acid. 

Solubility  and  Dissociation  of  Quinhydrone. — The  determination  of 
the  solubility  of  quinhy drone  is  complicated  by  the  fact  that  it  is  highly 
dissociated  in  aqueous  solution  into  its  two  components,  hydroquinone 
and  quinone. 

That  it  is  made  up  of  equimolecular  proportions  of  these  has  been 
shown  by  Liebermann,1  by  obtaining  the  maximum  yield  from  equi- 
molar  proportions  upon  mixing  aqueous  solutions  of  the  two;  by  Hesse2 
by  acetylation,  which  attacked  only  the  hydroquinone,  and  removing 
the  quinone  by  evaporation;  and  by  Nietzki,3  who  reduced  the  quinone 
with  SO2  and  titrated  back  the  excess  with  iodine.  The  same  fact  was 
also  confirmed  in  the  analytical  work  of  Valeur4  and  of  Luther  and  Leub- 
ner,4  and  similarly,  in  the  present  work,  in  which  it  was  further  verified 
by  mixing  together  equivalent  quantities  of  hydroquinone  and  quinone, 
in  solution,  and  determining  the  percentage  of  solids  in  the  filtrate  which 
checked  that  in  a  solution  saturated  with  quinhydrone  directly. 

In  Series  III5  the  hydrogen  ion  and  quinhydrone  concentrations  were 
kept  approximately  constant  and  the  hydroquinone  concentration  .was 
varied.  But  since  a  mobile  equilibrium  exists  in  solution  between  the  quinhy- 
drone and  its  dissociation  products,  we  have  also  a  varying  concentra- 
tion of  quinone  to  consider.  In  fact,  it  is  the  concentration  of  quinone 
rather  than  that  of  the  quinhydrone  in  which  we  are  primarily  interested. 
The  latter  does  not  enter  directly  into  the  reaction  whose  potential  we 
are  measuring,  but  its  unavoidable  presence  is  utilized  as  a  means  of 
regulating  the  concentrations  of  the  quinone  and  hydroquinone.  In 
order  to  know  what  these  are,  however,  we  must  know  not  only  the  true 
solubility  of  the  quinhydrone  but  its  equilibrium  constant  as  well.  By 
"true  solubility"  we  mean  the  concentration  of  dissolved  undissociated 

1  Ber.,  10,  1615  (1877). 

2  A.,  200,  248  (1877). 

3  Ber.,  10,  2000  (1877)  and  A.,  215,  130  (1882). 

4  Loc.  cit. 

6  See  page  22. 


24 

quinhy drone  in  equilibrium  with  the  solid.  The  quantity  of  quinhy  - 
drone  which  dissolves  in  saturating  the  solution  we  will  designate  as  the 
' '  apparent  solubility . ' ' 

In  Series  I  and  II  it  is  also  necessary  to  know  the  solubility,  etc.,  of  the 
quinhy  drone,  since  the  solubilities  do  not  remain  constant,  owing  to  the 
effect  of  the  hydrochloric  acid  upon  them. 

Luther  and  Leubner1  undertook  to  determine  the  true  solubility  and 
dissociation  constant  of  quinhydrone  in  water  by  the  method  used  by 
von  Behrend2  in  the  case  of  the  phenanthrene  picrates,  namely,  the  de- 
pression of  the  apparent  solubility  by  an  excess  of  one  of  the  dissociation 
products.  They  saturated  water  and  hydroquinone  solutions  of  known 
concentrations  with  quinhydrone  at  25°  C.  and  determined  the  total 
quinone  (combined  and  uncombined)  in  the  filtered  solution  by  Valeur's 
method,  which  is  perfectly  applicable,  owing  to  the  complete  dissocia- 
tion of  the  quinhydrone  as  the  quinone  is  removed  by  the  iodide.  This 
total  quinone  represents  (formula-weight  for  formula -weight)  the  total 
quinhydrone  which  has  dissolved  in  saturating  the  solution,  or  the  ap- 
parent solubility.  They  made  a  number  of  determinations  for  each  con- 
centration of  hydroquinone  taken,  the  averages  of  which,  as  given  in 
Table  IV,  they  used  in  their  calculations  as  follows: 

Let 

s  =  true  solubility  of  quinhydrone 


,  in  formula-weights  per  liter 
a  =  apparent  solubility  of  quinhydrone    ' 

b  =  excess  of  hydroquinone  added 

h  =  actual  concentration  of  hydroquinone  f  in  moles  per  liter 

q  =  actual  concentration  of  quinone 

Then 

total  hydroquinone  (combined  and  uncombined)  =  a  -}-  b  =  s-f-h  and  h  =  a  -f-  b —  s 
total  quinone  (as  determined  by  titration)  =  a  =  s  +  q,  and  q  =  a  —  s 

Let  K  =  *L*-Bf  and  P  =  Ks  =  h  X  q  =  (a  +  b  —  s)  (a  —  s)  (10) 

s 

If  the  "Law  of  Mass  Action"  held  perfectly,  and  s  remained  constant 
throughout  the  range  of  experiments,  K  would  be  a  constant,  namely, 
the  dissociation  constant.  If  the  experimental  precision  were  fine  enough, 
s  and  K  could  be  calculated  from  any  pair  of  determinations  by  means 
of  simultaneous  equations  of  the  form  of  equation  (10)  below. 

But  comparatively  slight  deviations  from  these  ideal  conditions  render 
this  method  of  calculation  inapplicable,  so  that  recourse  must  be  had 
to  a  method  of  trial  and  approximation  which  Luther  and  Leubner  car- 
ried out  in  the  following  way: 

By  trying  different  values  for  s  in  the  equation 

(a  +  b  — s)(a  — s)  =  P  (10) 

1  Loc.  cit. 

2  Zeit.  f.  physik.  Chem.,  15,  183  (1894). 


a. 

P. 

Deviation. 

_             P 

(Average.) 

a  -f  b. 

(s  »  0.0013.) 

from  mean. 

=  0.0013 

0.01827 

0.01827 

0.000288 

0.000009 

0.221 

0.01421 

O.O242I 

0.000296 

0.000001 

0.227 

O.OII50 

0.03150 

o  .  000308 

O.OOOOII 

0.236 

o  .  00664 

0.05664 

0.000296 

0.000001 

0.227 

they  found  the  value  for  s  for  which  P  showed  the  nearest  approach  to 
constancy,  that  is,  for  which  the  sum  of  the  deviations  of  P  from  its  mean 
value,  were  at  a  minimum.  This  value  for  s  they  found  to  be  0.0013, 
for  which  the  mean  value  of  P  was  0.000297,  and  the  sum  of  the  devia- 
tions O. 00002 2. 

TABLE  IV. 

a. 

b. 

O 

0.01 
0.02 
0.05 

We  repeated  the  work  of  Luther  and  Leubner,  extending  the  range 
up  to  the  saturation  point  for  hydroquinone,  in  water,  tenth-normal 
and  normal  hydrochloric  acid.  The  solutions,  with  excess  of  quinhy- 
drone,  were  placed  in  a  large  test  tube  fitted  with  a  spiral  mechanical 
stirrer  and  immersed  in  the  thermostat,  and  stirred  vigorously  and  con- 
tinuously. Samples  were  taken  about  every  fifteen  to  thirty  minutes, 
by  means  of  a  pipette  fitted  with  a  filter,  until  two  successive  titrations 
gave  the  s0me  result,  which  was  usually  the  case  with  the  first  two  sam- 
ples. In  a  number  of  cases  fresh  mixtures  were  made  up  and  tested  as 
a  check,  and  the  results  in  every  case  were  almost  identical  with  the 
original,  so  it  was  not  deemed  necessary  to  verify  all  the  solutions  in  this 
way,  the  regularity  of  the  results  and  the  parallelism  between  the  aqueous 
and  acid  solutions  also  serving  as  a  check. 

TABLE  V. 
Part  i.     Water.     No  hydrochloric  acid. 

Moles  per  liter. 


quinone 
d. 

j3  O 

+!<*.  V 
0  0  C 

0 

p 

X  10« 

for  values  of  s  at  heads  of 

columns. 

II 

III 

—  <  S 

H 

10 
tN 

CN 

o 

8 

00 

| 

| 

& 

<X!'C 

£ 

i 

8 

i 

§ 

8 

§ 

| 

§ 

b. 

a. 

b 

+  a. 

d 

d 

d 

d 

d 

o" 

o 

d 

0 

0.0178 

o 

.0178 

272 

274 

275 

282 

282 

283 

284 

284 

O.OI 

0.0135 

0.0235 

271 

273 

275 

28l 

282 

282 

283 

283 

0.02 

0.0106 

0 

.0306 

273 

274 

276 

284 

284 

285 

286 

286 

0.05 

0.00625 

o 

-05625 

273 

275  _ 

278 

290 

290 

291 

292 

293 

O.  I 

0.00374 

0 

•  10374 

250 

256 

263 

28l 

282 

284 

286 

287 

O.2 

0.00244 

0 

.20244 

229 

240 

250 

288 

290 

294 

298 

300 

0-3 

0.00189 

o 

.30189 

177 

193 

207 

265 

268 

274 

280 

283 

0.4 

O.OOI79 

0.40179 

197 

216 

237 

312 

316 

325 

333 

337 

0-5 

0.00172 

0 

.50172 

210, 

236 

260 

356 

361 

371 

38i 

385 

Sat'd 

0.00l8l5 

0.645+5 

332 

364 

395 

519 

526 

539 

55i 

558 

Sum  of 

deviations  of 

first 

four  values  of  P 

X 

I06 

from 

their  mean 

2 

4" 

J  J 

J  j 

I  I 

T   T 

Sum  of 

deviations 

of  first 

• 

J.   1 

seven 

from  their 

mean... 

1  86 

154 

127 

35 

32 

32 

32 

34 

Mean  of  first  seven  .  .                240     2  «;  * 

261 

282 

282 

28* 

287 

288 

K  for  s  = 
0.00098. 

0.289 
0.288 
0.291 

0.297 
0.290 
0.300 
0.280 
0.232 
0.378 
0.550 


26 


TABLE  V  (Continued). 
Part  2.     Tenth-Normal  Hydrochloric  Acid. 

P  X  106  for  values  of  s  at  heads  of  columns. 


b. 

a. 

b  +a. 

0.0013. 

0.0011. 

0.00103. 

0.00102. 

0.00101. 

is.  lor  s  = 
0.00102. 

0 

0.0173 

0.0173 

256 

262 

264 

265 

266 

0.260 

O.OI 

0.0131 

0.0231 

*      258 

264 

267 

267 

267 

0.262 

0.02 

O.OIO2 

o  .  0302 

257 

264 

267 

268 

268 

0.263 

0.05 

o  .  00593 

0-05593 

253 

265 

269 

269 

270 

0.264 

O.  I 

0.00363 

0.10363 

238 

260 

267 

268 

268 

0.263 

O.2 

0.00237 

0.20237 

215 

255 

270 

272 

274 

0.267 

0-3 

0.00190 

0.30190 

1  80 

241 

262 

265 

268 

0.260 

0.4 

0.00172 

0.40172 

1  68 

240 

276 

280 

284 

0.271 

0.5 

0.00170 

0.50170 

200 

300 

336 

341 

346 

0.329 

Sat'd 

0.00181 

0-633+5 

323 

450 

494 

500 

506 

0.490 

Mean  of 

first  seven  .  . 

.  .  237 

2SQ 

266 

268 

260 

Sum  of  deviations  of  first  seven  from 
their  mean 


156 


16 


Part  3. 


Normal  Hydrochloric  Acid. 

P  X  106  for  values  of  s  at  heads  of  columns. 


b. 

a. 

b  +  a. 

0.0012. 

0.0010. 

0.00088.  < 

0.00087. 

0.00086. 

iv.  tor  s  = 
0.00087. 

O 

O.OI62 

O.OI62 

225 

231 

235 

235 

236 

0.270 

O.OI 

O.OIlS 

0.0218 

219 

225 

229 

229 

230 

0.263 

0.02 

0.0091 

0.0291 

220 

228 

232 

232 

232 

0.267 

0.05 

0.0052 

0.0552 

216 

228 

235 

235 

236 

0.271 

O.I 

0.0031 

0.1031 

194 

215 

227 

228 

229 

O.262 

0.2 

0.00202 

O  .  2O2O2 

I65 

206 

230 

232 

234 

0.267 

0-3 

0.00164 

0.30164 

134 

193 

229 

232 

235 

0.267 

0-4 

O.OOI52 

0.40152 

128 

208 

256 

260 

264 

0.299 

Sat'd 
Mean  of 

0.00159 

first  seven.. 

0-494+s 

193 
196 

292 
218 

351 
231 

356 
232 

361 
233 

0.409 

Sum  of  deviations  of  first  seven  from 

their  mean 191  80  18  13  17 

Attention  is  called  to  the  peculiar  fact  that  in  ail  three  cases  at  the  saturation 
point  of  hydroquinone,  the  apparent  solubility  of  quinhydrone  actually  increases  in- 
stead of  decreasing  with  increasing  excess  of  hydroquinone. 

The  results  are  given  in  Table  V.  On  comparing  our  results  for  water 
solutions  with  those  of  Luther  and  Leubner,  we  note  the  following:  In 
the  first  place,  our  results  are  lower  than  theirs  by  about  5%  of  the  ap- 
parent solubilities  themselves.  We  do  not  know  the  reason  for  this.  In 
the  second  place,  the  value  of  s,  which  gives  the  minimum  sum  of  devia- 
tions of  P  from  its  mean,  for  the  range  comprising  our  first  four  solutions, 
which  was  as  far  as  their  investigation  was  carried,  is  0.00125,  which 
checks  their  value  0.0013  by  less  than  5%.  But  the  minimum  sum 
of  our  deviations  for  this  range  is  only  o .  000002  whereas  theirs  is  o .  000022, 
eleven  times  as  much.  From  this  it  appears  that  we  have  been  able  to 
obtain  a  much  greater  degree  of  precision. 


27 

If,  however,  we  consider  the  entire  range  up  to  the  saturation  point 
of  hydroquinone,  we  see  in  the  column  of  s  =  0.00125  a  decided  but  con- 
tinuous decrease  in  P  reaching  a  minimum  when  b  =  0.3,  followed  by 
a  continuous  increase  becoming  suddenly  abrupt  at  the  saturation  point. 
It  was  to  verify  these  seemingly  abnormal  results  that  determinations 
were  repeated  which  satisfied  us  that  they  were  not  due  to  experimental 
error.  If,  now,  we  try  smaller  values  of  s,  we  find  that  our  first  seven 
values  of  P  become  more  and  more  uniform,  givdng  a  minimum  devia- 
tion sum  of  0.00032  for  values  of  s  between  o.ooioo  and  0.00096.  We 
therefore  select  the  middle  value  0.00098  as  the  most  probable  value 
of  s  in  water,  according  to  our  experiments.  This  involves  the  assump- 
tion as  a  working  hypothesis,  in  preference  to  the  only  other  and  much 
less  probable  alternative,  viz.,  that  they  vary  inversely  to  each  other, 
thus  giving  a  constant  product,  that  both  s  and  K  remain  approximately 
constant  for  values  of  b  up  to  0.3.  Above  this  value,  one  or  both  must 
increase.  A  clue  as  to  which  one  seems  to  be  given  to  us  by  our  potential 
measurements,  as  will  be  shown  in  the  next  section. 

Perfectly  parallel  results,  it  will  be  noticed,  were  obtained  with  tenth- 
normal  and  normal  hydrochloric  acid  solutions.  The  solubilities  and 
dissociation  constants  of  quinhy drone  are  recapitulated  in  Table  VI. 
The  basis  of  the  last  column  is  to  be  found  in  the  next  section. 


TABLE  VI. 

Moles  per  liter. 

Dissociation  constant. 

Solvent. 
Water  

Apparent 
solubility. 

.  .       o  0178 

True 
solubility. 

0.00098 
O.OOIO2 
0.00087 

Normal  . 
(Av.  1st  7.) 

0.289 
0.263 
o.  267 

Saturated 
hydroquinone. 

0.550 
0.490 
O.  AOQ 

N/io  HC1  

O  OI7"? 

#HC1.. 

O.OI62 

III.    THEORETICAL   SIGNIFICANCE   OF   THE   POTENTIALS. 

If  the  potentials  observed  measure  mainly  the  tendency  to  take  place 
of  a  reaction  represented  by  the  following  equation: 

C6H402  +  2H+  +  2e  ^  C6H4(OH)2  (n) 

then,    theoretically,  we  should  find  the  following  relation  to    hold  ap- 
proximately • 

RT  r       q  X  [H+]2 


" 


0.0298  £log  q  Xh[H+l2  -  _  log  Kol  volts  for  25°  C.         (12) 
or,  as  this  form  of  equation  is  usually  written, 

7T    =    0.0298  log  q    X    [H  +  ]2    +    TTo     = 


0.0596  log    [H  +  ]    +   0.0298  log       ;          (I3) 

n 


28 

in  which  TT  is  the  observed  potential  for  the  concentrations  q,  [H+]  and 

q  X  [H+l2 
h,  K0  is  the  value  of  ~ ^ when  TT  =  o,  i.  e.,  at  "equilibrium,"  and 

q  x  [H+l2 
7T0  is  the  value  of  IT  when ^ =  i,  for  then  the  first  term  0.0298 

log  3  __^ —    -  =  o,  and  TTO  =  — 0.0298  log  K0. 

But  so  far,  we  have  data  for  h  and  q  at  our  disposal  only  in  the  cases 
of  mixtures  F,  G,  H,  I  and  J. 

In  the  last  form  of  the  above  equation  (Equation  (13)),  the  first  term 
TTO  is,  by  definition,  a  constant,  and,  in  the  tenth-normal  hydrochloric 
acid  solutions,  the  second  term  0.0596  log  [H+]  may  be  assumed  to  be 
nearly  a  constant.  Therefore,  the  differences  (next  to  the  last  column, 


TABUJ  VII. 

Difference. 


b.     q«=a —  s.    h  =  a+b  —  s.     q/h.     0.0298  log  q/h.     Calculated.     Found. 

J o   0.01628  0.01628  i. ooo      o    j 

I o.oi  0.01208  0.02208  0.547   —0.00781! 

TT  O.OO74I      O.OO65 

H 0.02     0.00918     0.02918     0.315      — 0.01495} 

«  0.0162      0.0155 

G 0.05     0.00491     0.05491     0.0895     — 0.0312  j  v* 

F o.i       0.00261     0.10261     0.0254    —0.0475  ' 

The  figures  in  this  table  are  derived,  as  indicated,  from  those  of  Table  V,  Part  2, 
which  see.     The  column  "b"  is  common  to  both  tables. 

Table  VII)  between  the  values  of  0.0298  log  q/h  may  be  taken  as  ap- 
proximating the  theoretical  potential  differences  themselves.  That  is, 
for  this  series,  theoretically, 

TTi  —  7T2  =  0.0298  log  (q/h)i  —  0.0298  log  (q/h)2.  (14) 

The  observed  potential  differences  are  represented  by  the  distances 
between  the  lines  in  Chart  I.  Each  pair,  consisting  of  a  light  and  a  heavy 
line,  represents  the  duplicate  cells  of  the  mixture  indicated  by  the  letter 
at  the  right-hand  end  of  the  pair  of  lines.  The  heavy  lines  connect  the 
points  representing  the  readings  of  cell  No.  i  and  the  light  lines  those  of 
cell  No.  2  of  the  mixtures  in  question.1  The  distances  between  the  various 
pairs  of  lines,  it  will  be  seen,  correspond  quite  distinctly  to  the  theoretical 
differences2  which  are  indicated  by  the  distances  between  the  horizontal 
dotted  lines  in  the  chart.  The  manner  in  which  the  absolute  values  repre- 
sented by  these  dotted  lines  were  obtained,  will  appear  a  little  later  in 
this  discussion. 

Before  we  can  arrive  at  a  definite  single  figure  to  be  taken  from  our 
experimental  data  as  the  true  representative  experimental  potential  value 

1  Except  in  the  case  of  B  in  which  they  are  No.  3  and  No.  4,  respectively. 

2  Table  VII,  "Differences  calculated,"  TTI  —  TTZ,  Eq.  (14). 


CHART  I -SERIES  3 


.94 


.92 


20  40  60  80  100  120 


160  ISO 


(for  purposes  of  numerical  comparison  and  calculation)  for  each  mixture, 
several  things  must  be  taken  into  account.  In  the  first  place,  we  notice 
a  well-defined  and  quite  regular  sloping  off  of  the  potentials,  which  be- 
comes more  marked  as  the  proportion  of  quinone  increases,  indicating 
one  or  more  side  reactions  involving  quinone,  and  associated  very  likely 
with  the  increasing  brown  color  which  solutions  of  quinone  acquire  on 
standing,  and  consistent  with  the  general  instability  or  reactivity  of  qui- 
none. But  what  we  want  is  the  potential  of  the  mixture  as  we  have 
made  it  up,  before  its  original  composition  has  been  altered  by  any  side 
reactions.  The  natural  solution  of  the  difficulty  would  be  to  take  the 
average  initial  or  first  day's  readings,  but  it  will  be  noticed  that  the  read- 
ings taken  during  the  first  twenty-four  hours  are  usually  quite  erratic 
and  non-reproducible,  indicating  the  influence  of  accidental  external 
factors  before  they  have  come  to  equilibrium  with  the  principal  system. 
To  obtain  the  true  initial  potentials,  therefore,  we  are  compelled  to  re- 
sort to  a  sort  of  graphical  extrapolation,  by  inspection  instead  of  calcula- 
tion. 

Without  laying  down  any  hard  and  fast  rule,  we  have  selected  as  the 
true  initial  potential  for  each  of  the  mixtures  in  question  that  value  which, 
upon  inspection,  ignoring  the  irregularities  of  the  first  day  or' two,  seems 
to  be  most  consistent  with  the  contours  of  the  curves  for  that  mixture 
and  with  those  of  the  neighboring  curves.  The  values  so  selected  are 
marked  on  the  chart  by  the  short  lines  extending  to  the  left  from  the  left 


30 

border  of  the  diagram.  The  fairness  of  this  method  of  approximation 
and  its  precision  of  half  a  millivolt,  which  is  sufficient  for  the  purpose, 
may  be  seen  on  inspection  of  the  chart.  The  method  is  not  as  crude  as 
it  might  seem,  at  first  glance,  and  is  the  only  one  suitable  for  the  purpose. 

The  values,  "TT,"  in  Table  IX1  are  these  values,  and  the  "Differences 
found,"  in  the  last  column  of  Table  VII,  are  the  differences  between  them. 
It  will  be  noticed  that  the  observed  differences  ("found")  do  not  differ 
greatly  from  the  theoretical  differences  ("calculated"). 

Thus  our  experimental  results  are  found  to  be  in  rough  agreement 
with  the  van't  Hoff  equation,  in  the  form 

TT  =  7T0+  0.0298  log  (q/h)  +  0.0596  log  [H+]  (13) 

in  the  case  of  the  potentials  of  those  mixtures  for  which  we  have  the  neces- 
sary concentration  data.  But  in  the  case  of  the  solutions  saturated  with 
hydroquinone  (mixtures  A,  B  and  C),  this  concentration  data  is  lacking, 
owing  to  the  fact  already  pointed  out  that  the  constancy  of  both  s  and  K, 
upon  which  the  method  of  calculation  depends,  does  not  continue  up  to 
the  saturation  point  of  the  solution  with  respect  to  hydroquinone,  so 
that  the  method  is  inapplicable.  But  if  we  make  the  tentative  assumption 
that  the  above  expression  (Eq.  (13))  holds  also  in  this  case,  we  have  a 
means  of  calculating  the  concentration  ratio  (q/h),  and  from  it  the  con- 
centration of  quinone  (q),  the  solubility  of  quinhy drone  (s),  and  the  value 
of  the  dissociation  constant  (K).  We  are  then  in  a  position  to  apply 
the  above  equation  to  the  study  of  the  effect  of  varying  acidity  in  the 
case  of  the  stable  potentials  of  the  saturated  hydroquinone  solutions 
(Series  I). 

Inspection  of  the  chart  brings  out  the  fact  that  mixture  F  gives  the 
most  stable  potential  of  the  group  already  considered.  So  we  select  it 
as  the  basis  for  our  calculations,  taking  for  TTJ?,  the  initial  reading  (o .  8695) 
obtained  as  explained  above,  which  is  also  the  average  of  the  readings 
up  to  the  point  at  which  the  falling  off  commences.  For  TTB  we  have 
the  average  value  0.8300. 

We  have,  then, 

TTF  —  TTB  =  0.0298  log  (q/h)F  —  0.0298  log  (q/h)B  (14) 

or,  0.8695  —  0.8300  =  —0.0475*  —  0.0298  log  (q/h)B. 

Solving          (q/h)B  =  -log-{0-039o5+9°80475)  =  o.ooI202; 

hB  =  o.633,2soqB  =  0.001202  X  0.633  =  0.00076; 

1  Page  33. 

*  See  Table  VII. 

2  See  Table  II. 


SB  =  a  —  q  =  o.ooiSi1  —  0.00076  =  0.00105; 
pB  =  q  x  h  =  0.00076  X  0.633  =  0.000481; 

P        0.000481 

K    =  -  =  -  -  =  0.457. 

s        0.00105 

Now  the  difference  between  the  value  (0.00105)  obtained  for  s,  in 
saturated  hydroquinone  solution  by  this  method,  and  the  value  (0.00102) 
obtained  in  the  last  section2  for  that  range  over  which  s  and  K  appear 
to  be  approximately  constant,  is  well  within  the  limits  of  certainty  of 
the  method  by  which  the  value  (0.00105)  was  obtained.  This  becomes 
e vide  at  when  we  repeat  the  above  calculation,  using  each  of  the  other 
mixtures  of  the  series  as  the  basis,  in  place  of  mixture  F. 

Bans.  SB. 

F 0.00105 

G 0.00105 

H . O.OOIOI 

I o . 00094 

J o . 00092 

The  average  of  these  values  is  only  0.00099,  but  the  higher  values  are 
based  on  the  mixtures  giving  the  more  stable  and  definite  potentials 
and  therefore  carry  more  weight.  All  that  we  are  really  justified  in  say- 
ing, then,  is  that  apparently  the  true  solubility  (s)  of  quinhydrone,  re- 
mains about  constant  and  it  is  only  its  dissociation  "constant"  which  in- 
creases suddenly  as  the  saturation  point  for  hydroquinone  is  approached. 

For  our  purpose,  therefore,  we  consider  excess  of  hydroquinone  to  be 
without  influence  on  the  true  solubility  of  quinhydrone.  We  have  re- 
calculated our  data  for  Mixture  B  on  this  basis  (s  equals 0.00102  instead 
of  0.00105).  The  results  are  recorded  in  Table  VIII.  We  did  not  run 
a  special  series  of  solubility  experiments  for  N/ioo  HC1  because  there  is 
so  little  difference  between  A7/io  HC1  and  water  that  interpolation  would 
show  a  negligible  difference  between  water  and  N/ioo  HC1,  and  we  thought 
that  data  for  water  as  solvent  would  be  of  more  general  value  than  data 
for  N/ioo  HC1.  We  also  wanted  to  compare  our  results  with  those  of 
Luther  and  Leubner,  who  worked  only  with  solutions  in  pure  water. 
Table  VIII  is  based  on  the  equation 

TA  —  TB  =  (0.0298  log  (q/h)A.+  0.0596  log  [H+]A)  - 

(0.0298  log  (q/h)B  +  0.0596  log  [H+]B).         (15) 

TO  cancels  out.  The  values  under  "IT  (observed)"  are  the  observed  aver- 
age values  (taken  from  Chart  II,  in  which  all  of  the  cells  of  the  mixtures 
indicated  are  plotted  from  data  of  Table  I,  as  in  Chart  I)  of  the  poten- 
tials, during  their  period  of  stability. 

iSee  Table  V,  Part  2. 
2  See  Table  V. 


TABLE  VIII. 


Mix-         q  = 
ture.       a  —  s. 


Calcu- 
lated. 


h.         [H+].         (0.0298  log  q/h)  + (0.0596  log  [H+]. 

A  0.00072  0.4940.790  (—0.0845) +  (-0.00609)  =—0.0906 

B  0.00079  0.633  0.0921  ( — 0.0867)  +  (-0.0618  )= — 0.1485 

C  0.000835  0.645  0.00971  ( — o.o86i)  +  (-o.i20o  )= — 0.2061  J°-°576 


0.0579 


Ob-         Ob- 
served, served. 

°-883}  0.053 
0.830 
0.770!  0-060 


In  like  manner,  TTA  —  TTC  =  0.1155  (calculated)  and  o.  113  (observed). 
Thus  the  observed  potentials  are  found  roughly  to  be  in  agreement  with 
the  theory  in  the  case  of  varying  hydrogen-ion  concentration  also 


CHART  E- SERIES  1 8.2 


0  50 


300 


350 


400  450 


The  values  under  [H+]  were  obtained  as  follows:  Bray  and  Hunt,1  by 
the  conductivity  method,  found  for  the  degree  of  ionization  (a)  of  hydro- 
chloric acid  at  25°,  92.1%  in  tenth-normal  solution,  and  97.1%  in  hun- 
dredth-normal solution.  But  no  direct  data  could  be  found  in  the  litera- 
ture for  normal  hydrochloric  acid  at  25°.  Kohlrausch,  however,2  gives 
the  following  for  the  equivalent  conductivity  (A)  at  18°: 

Normality.  A. 

0.01  370 

O.I  351 

i.o  301 


From  these  figures  we  find  A.i/A.oi  =  0.948  at  18° 
and  Hunt's  data  we  find  that  a.\/a.n  =  0.948  at  25°. 

1  /.  A.  C.  S.,  33,  781  (1911)- 

2  Landolt-Bornstein  Tab.,  2410.,  p.  1104. 


while  from  Bray 
Moreover,  accord- 


33 

ing  to  Kohlrauscb,1  the  temperature  coefficient  of  conductivity  of  HC1 
solutions  varies  less  and  less  with  increasing  concentration,  approaching 
constancy  as  normal  concentration  is  approached.  Thus: 

Concentration o.ooi  o.oi  o.  i  o.s 

Coefficient 0.0163  0.0158  0.0153  0.0152 

(The  coefficient  for  normal  hydrochloric  acid  is  not  given.) 
In  the  absence  of  definite  data,  therefore,  we  assume  that 

(*1\      orf  — )       =  (— }       =  0.858  (from  Kohlrausch's  figures) 

\A.i/25°          \«.l/25°  \A.i/i8° 

ai  =  a.i(  — )       =0.921X0.858  =  0.790. 

\A.i/i8° 

3o  we  take,  as  an  approximation,  the  value  0.790  for  the  hydrogen  ion 
[Concentration  in  our  normal  hydrochloric  acid  solution  (A). 

The  method  used  by  the  leading  workers  in  this  field  in  inorganic  chem- 
istry2 for  comparing  observed  potentials  with  the  theory,  was  to  calcu- 
late, for  each  observed  potential,  by  means  of  an  equation  similar  in  form 
to  the  one  we  have  been  using  (Equation  (12)),  the  value  of  the  expression 
— RT/nF  loge  K0  (designated  by  Peters  and  Friedenhagen  as  A  and  by 
Luther  and  in  the  present  work  as  TTO),  the  degree  of  constancy  of  these 
values  measuring  the  agreement  with  the  theory.  Applying  this  method 
to  our  own  results,  we  find  the  following : 

7T0  =  TT  —  0.0596  log  [H+]  —  0.0298  log  q/h  (16) 


TABLE 

IX. 

F. 

G. 

H. 

I. 

J. 

B. 

C. 

A. 

-0.0596  log  [H  +  ] 
r,  

.      0. 
.      0. 

0618 
8695 

0. 
0. 

0618 
8860 

0.0618 
0.9015 

0.0618 
0.9075 

0 
0, 

,0618 
9150 

0.0618 
0.8300 

0.1200 
0.7700 

0.0061 
0.8830 

-0.0298  log  q/h.. 

.     0. 

0475 

0. 

0312 

0.0149 

0.0078 

0, 

0000 

0.0867 

0.0861 

0.0845 

ro  .  . 

.      0. 

9788 

0 

,9790 

0.9782 

0.9771 

0, 

,9768 

0.9785 

0.9761 

0.9736 

The  deviations  (Table  IX)  are  of  the  respective  values  of  ir0  from  their 
mean,  which  is  0.9773.  The  average  deviation,  or  numerical  mean  of 
these  deviations  is  0.0014  volt,  or  i  .4  millivolts.  This  compares  favor- 
ably with  the  closest  agreement  that  has  been  obtained  in  parallel  work 
in  inorganic  chemistry,  as  can  be  seen  from  the  following  summary: 

1  Loc.  cit.,  244  b.  p.  1115. 

2  See  articles  by  Peters,  Friedenhagen  and  Luther,  already  cited,  p.  10. 


34 


System. 

Ferrous /ferric 

Ferrous /ferric 

Ferrocyanide/ferricyanide 

Manganate/permanganate 

Uranous/uranyl 

Iodine /iodate 

B  romine  /br  ornate 

Bromide/bromine 

Iodide/iodine 

Hy  droquinone  /quinone 


nvestigators. 

Peters  (Tab.  i) 
Peters  (Tab.  2) 
Friedenhagen 
Friedenhagen 
Luther  &  Michie 
Luther  &  Sammet 
Luther  &  Sammet 
Luther  &  Sammet 
Luther  &  Sammet 
This  work 


Average 
deviation. 
Millivolts. 

39 
2-7 
i.o 
4-o 

I  .2 

0.8 

I.O 

0.5 

I  .O 

14 


From  0.0298  log  Kc 

,-33 


9773>  we  can  calculate  the  value  of  Kc 
which  is  i  .6  X  10 

We  may  also  use  the  mean  value  TTO   =  o.  9773  as  a  basis  for  calculating 
theoretical  potentials  of  the  mixtures  under  consideration  as  follows: 


TABLE  X. 

Mixture. 

7TO               + 

.0.0298  log  l(q/h)   [n+]2  1  = 

TT  (calculated). 

TT  (observed). 

A 

0.9773 

—  o  .  0906 

0.8867 

0.8830 

B 

0-9773 

—0.1485 

0.8288 

o  .  8300 

c 

0-9773 

—  0.2061 

0.7712 

0.7700 

F 

0-9773 

—  0.1093 

0.8680 

0.8695 

G 

0.9773 

—  0.0930 

0.8843 

0  .  8860 

H 

0-9773 

—  0.0767 

0.9006 

0.9015 

I 

0-9773 

—  o  .  0696 

0.9077 

0.9075 

'J 

0-9773 

—  0.0618 

0.9155 

0.9150 

The  relation  of  the  theoretically  calculated  potentials  to  the  observed 
can  best  be  seen  by  graphical  means.  They  are  represented  on  the  charts 
by  the  straight  horizontal  dotted  lines. 

In  concluding  this  section,  attention  is  called  to  the  following  points: 
It  was  not  found  possible  to  determine  the  concentrations  directly  or  to 
compute  them  with  mathematical  certainty  from  the  experimental  data. 
Instead,  the  values  used  have  been  derived  from  the  experimental  data 
by  a  method  based  upon  probability.  Therefore,  it  must  be  admitted 
that  the  method  by  which  they  were  obtained,  alone,  does  not  justify 
us  in  presenting  them  as  anything  more  than  probable  values.  But  this 
method  has  been  clearly  and  definitely  recorded,  so  that  the  doubt  exists 
only  as  to  their  significance.  The  fact  has  been  established  that  the 
values  so  obtained,  when  taken  as  the  actual  concentrations,  fit  the  elec- 
trochemical theory  to  the  degree  of  precision  shown.  This  coincidence 
is  so  remarkable  that  it  greatly  strengthens  the  probability  that  these 
values  approximate  the  actual  concentrations,  or  at  least  the  effective 
concentrations. 


PART  H. 


I.    POTENTIALS  OF  SOLUTIONS  SATURATED  WITH  QUINONE. 

Only  two  mixtures  saturated  with  quinone  were  tried,  viz.,  D  and  E, 
already  described.  In  Table  I  the  measurements  on  two  cells  of  each 
of  these  mixtures  are  given.  The  two  E  cells  are  the  only  ones  that  were 
tried  of  this  mixture.  But,  in  the  case  of  D,  the  two  given  are  the  best 
specimens  out  of  five.  The  first  three  are  given  in  Table  XL 

TABLE  XL 


Days. 

Cell  No.  1. 

0 

I  .  0064 

I 

0.9903 

2 

o  .  9860 

3 

0.9855 

4 

0.9835 

5 

0.9835 

6 

o  .  9805 

7 

0.9780 

8 

o  .  9740 

9 

0.9705 

10 

o  .  9648 

ii 

0.9565 

12 

0.9487 

13 

o  .  9390 

14 

0.9345 

15 

0.9285 

16 

0.9235 

17 

0.9200 

18 

0.9180 

19 

0.9156 

20 

0.9150 

21 

0.9138 

22 

0.9125 

23 

0.9130 

24 

0.9130 

25 

0.9125 

26 

0.9115 

27 

0.9110 

28 

0.9110 

29 

0.9104 

30 

0.9103 

31 

0.9100 

32 

o  .  9094 

33 

o  9092 

Cell  No.  2. 

Cell  No.  3. 

Days. 

Cell  No.  1.  (cont'd) 

0.9901 

o  .  9842 

34 

0.9077 

o  -  9749 

0.9761 

35 

0.9I502 

o  .  9668 

o  .  962  i 

36 

0.9597 

0.9505 

37 

38 

0.9389 

0.9277 

39 

0.9230 

o  .  9076 

40 

0.9120 

0.9147 

0.8937 

41 

0.9107 

0.8976 

0.8661 

42 

0.8888 

0.8310 

43 

0.8723 

0.7885 

44 

o  .  9082 

45 

0-8573 

0-7559 

46 

o.9io63 

0.8474 

0.7423 

47 

o  9090 

0.8322 

0.7238 

48 

o  .  9070 

0.8316 

0.7175 

49 

o  9049 

0.8275 

o  .  662  i 

50 

0.8295 

o  .  6248 

51 

o  .  9088 

52 

0.9081 

0.8180 

0.5903 

53 

0.9083 

54 

o  .  9068 

0-7935 

0.5747 

55 

0.9063 

56 

o  .  9062 

0.7849 

0.5719 

57 

0.7715 

0.5724 

58 

59 

0.9041 

o  7605 

0.5764 

60 

o  .  9048 

61 

0.9041 

0-7499 

0.5736 

0.7489 

o.57281-o.955 

8 

o  -  7493 

0-9375 

0.7452 

0.9227 

0.7410 

0.9108 

o  .  7409 

1  At  this  point  it  was  found  that  practically  all  of  the  undissolved  solids  had  dis- 
appeared.    So  the  cell  was  opened  and  more  quinone  and  quinhydrone  added  and 
nitrogen  passed  for  a  few  hours.     The  two  readings  on  that  day  were  taken  before  and 
after  adding  the  solids. 

2  On  this  morning  it  was  found  that  the  vibrator  on  the  relay  of  the  temperature 


36 

The  first  cell,  it  will  be  seen,  was  measured  over  a  period  of  two  months, 
during  which  time  it  fell  steadily,  except  for  occasional  periods  of  apparent 
equilibrium  lasting  a  few  days.  In  all  our  cells  the  electrode  was  im- 
bedded in  an  excess  of  the  solids  with  which  the  solution  was  saturated. 
In  the  case  of  this  cell,  it  was  noticed  that  the  yellow  quinone  particles 
soon  became  dark  green,  and  this  change  was  naturally  associated  with 
the  falling  potential.  It  was  thought  that  the  dark  product  might  be, 
partly  at  least,  quinhy drone.  To  test  this  idea  the  second  cell  was  made 
up  without  the  addition  of  quinhydrone,  the  idea  being  that  if  the  quinone 
produced  its  own  quinhydrone  the  solution  would  eventually  become 
saturated  with  this  substance,  in  which  event  the  potential,  though 
originally  higher,  would,  other  conditions  being  equal,  become  about  the 
same  as  that  of  cell  No.  i.  Cell  No.  3  was  made  up  similarly  to  cell  No. 
i,  but  with  less  excess  solids,  and  started  simultaneously  with  Cell  No.  2. 
It  will  be  noticed  that  the  potential  of  No.  2  is  from  the  start  intermediate 
between  that  of  No.  i  and  No.  3.  The  dark  solid  was  observed  in  Cell 
No.  2  on  the  second  day,  so  its  formation  begins  early.  No  definite  con- 
clusions on  this  point  could  be  drawn,  owing  to  unknown  conditions, 
however,  and  the  matter  was  not  considered  of  sufficient  importance  to 
the  main  investigation  to  be  pursued  further  at  this  time.  On  comparing 
Cells  No.  i  and  No.  3  we  see  that  evidently  the  excess  of  solids  has  a  great 
deal  to  do  with  the  stability  of  the  potentials.  This  is  strikingly  borne 
out  by  the  great  increase  in  potential  to  nearly  the  original  value,  on 
adding  fresh  solids  to  No.  3  on  the  2 9th  day.  The  sudden  increase  was 
not  due  to  exposure  to  the  air  on  opening  the  cells,  as  it  did  not  begin 
until  after  the  solids  were  added.  The  fourth  and  fifth  cells  were  there- 
fore made  up  with  a  much  greater  excess  of  solids,  as  previously  men- 
tioned, expecially  quinone,  as  were  also  the  tenth-normal  hydrochloric 
acid  cells  (Mixture  E),  giving  the  decidedly  more  stable  potentials  already 
recorded  in  Table  I. 

By  means  of  the  equation 

q[H+2l 

7T    =    0.0298  log  ^-L- -"    -f  TTo 

n 

and  the  data  of  the  preceding  sections,  we  can  calculate  a  theoretical 
potential  for  the  mixture  B. 

We  have  (assuming  K  and  s  to  be  unaffected  by  an  excess  of  quinone 
to  saturation  (0.1275  molar) 

regulator  had  stuck  during  the  night  permitting  the  temperature  of  the  thermostat 
to  rise  to  39°  C.  The  bath  was  immediately  cooled  to  25°  again,  but  this  temporary 
heating  seems  to  have  had  a  permanent  elevating  effect  on  the  potential. 

3  Nitrogen  was  passed  again  for  a  few  hours  to  see  what  effect  the  stirring  and  re- 
moval of  gases  might  have  on  the  potential.  It  seems  to  have  had  a  slight  elevating 
effect. 


37 

h  X  q  =  0.000268  (Table  V,  Part  2);  h  =  -  -   =  0.002102; 

[H+]  =  0.0921  (Table  VIII);  TTO  =  0.9773  (Tables  IX  and  X); 

o.  1275  X  (o.092i)2 
*  =  0.9773  +  0.0298  log  -  VV  =  0-9687. 


This  value,  it  will  be  seen,  is  higher  than  even  the  initial  values  (0.9517 
and  0.9524)  observed  for  this  mixture,  and  considerably  higher  than 
the  apparent  equilibrium  (horizontal  part  of  the  potential-time  curves) 
values.  We  have  already  seen,  in  Series  III  as  a  whole,  that  there  is  a 
distinct  falling  short  of  the  theory,  which  becomes  more  and  more  marked 
as  the  proportion  of  quinone  is  increased.  Furthermore,  the  increasing 
instability  of  the  potentials  parallels  this  deviation  from  the  theory. 
These  facts  point  to  the  conclusion  that  the  falling  off  from  the  theory  is 
connected  in  some  way  with  the  instability  of  quinone  rather  than  to  the 
failure  of  the  theory.  Considering  the  length  of  time  required  for  the  ap- 
parent equilibrium  to  be  reached,  in  mixtures  D  and  E,  in  connection  with 
the  instability  of  quinone,  it  seems  most  probable  that  the  initial  poten- 
tials are  closer  to  the  true  potential  of  the  system  as  it  was  made  up  origin- 
ally than  the  constant  potentials  reached  later  are.  Yellow  quinone 
particles  could  still  be  distinguished  in  the  solid  mass  around  the  elec- 
trode when  cells  No.  4  and  No.  5,  of  Mixture  D,  and  the  cells  of  Mixture  E, 
were  finally  removed,  and  the  odor  of  quinone  was  marked.  But  even 
though  the  solution  remained  saturated,  which  is  not  certain,  products 
may  have  been  formed  which  affected  either  the  potential  directly  or  the 
hydrogen-ion  concentration.  It  is  evident  that  something  affected  the 
potential. 

Hydrochloric  acid  solutions  of  quinone  and  also  of  quinhydrone  soon 
acquire  a  deep  rose  color  on  standing,  which  is  not  extracted  by  ether. 
This  color  is  noticeable  even  in  Vioo  normal  hydrochloric  acid.  Accord- 
ing to  Beilstein,  dilute  hydrochloric  acid  is  without  action  on  quinone. 
When  solid  quinone  is  treated  with  concentrated  hydrochloric  acid,  it 
first  becomes  dark  and  then  dissolves  to  a  colorless  solution.  The  final 
product  is  monochlorhydroquinone.1  That  this  reaction  is  not  confined 
to  the  concentrated  conditions  described  in  the  literature  was  easily  shown 
by  a  simple  rough  qualitative  experiment.  A  series  of  portions  of  quinone 
were  treated,  respectively,  with  hydrochloric  acid  in  concentrations 
varying  from  concentrated  (12-13  normal)  to  normal.  With  the  con- 
centrated acid  the  reaction  took  place  immediately.  The  others  followed 
in  succession  in  the  order  of  their  concentrations.  After  twenty-four 
hours,  the  first  four  (including  7  .  5  normal)  had  become  colorless  ;  the  next 
one  (6  normal)  was  well  on  the  road  to  colorless;  the  others  were  successively 
'Wohler,  A.,  51,  155. 


38 

darker,  the  last  three  (3.5,  2  and  i  normal)  being  just  at  the  darkest 
intermediate  stage.  After  two  days  more  no  further  changes  had  taken 
place.  The  complete  reaction  was  also  obtained  by  adding  concentrated 
hydrochloric  acid  to  an  aqueous  solution  of  quinone.  Now  this  reaction 
involves  two  changes:  (i)  reduction  to  hydroquinone,  and  (2)  chlorina- 
tion  of  the  hydroquinone,  as  represented  in  the  two  following  equations : 

(1)  C6H402  +  2HC1  — >  C6H4(OH)2  +  C12 

(2)  C6H4(OH)2  +  C12  — >  C6H3(OH)2C1  +  HC1 

That  the  reaction  does  probably  go  through  these  two  steps,  one  at  a 
time,  and  that  the  dark  intermediate  product  is  probably  mainly  quin- 
hydrone,  was  demonstrated  in  the  following  way:  Concentrated  hydro- 
chloric acid  was  added  to  a  saturated  aqueous  solution  of  quinone,  drop 
by  drop,  until  darkening  just  began.  The  darkening  continued  without 
further  addition  of  acid,  and  the  characteristic  bronze-green  product 
soon  separated  out.  It  was  recrystallized  from  glacial  acetic  acid.  The 
product  had  the  appearance  characteristic  of  quinhy drone.  The  air 
dried  crystals  gave  a  titration,  by  the  Valeur  method,  corresponding  to 
92 . 8%  quinhy  drone.  No  further  confirmation  was  made  at  this  time. 
Here  we  have,  evidently,  a  simple  case  of  reduction  by  hydrochloric  acid. 
The  rose  color,  mentioned  above,  was  also  noticed  as  an  intermediate 
stage  in  the  reaction.  The  whole  reaction  should  form  an  interesting  sub- 
ject for  a  special  investigation,  expecially  because  of  the  fact  that  passing 
chlorine  gas  into  a  solution  of  hydroquinone  oxidizes  it  through  quin- 
hy drone  completely  to  quinone,  instead  of  chlorinating  it;  that  is,  it  re- 
verses the  first  of  the  above  reactions  instead  of  causing  the  second  to 
take  place. 

From  the  above  it  is  evident  that  quinone  does  react  with  hydrochloric 
acid,  even  in  dilute  solution,  which  means  a  lowering  of  hydrogen-ion 
concentration  and  corresponding  lowering  of  potential.  This  may  ac- 
count in  part  for  the  falling  quinone  potentials  and  perhaps  also  for  their 
low  initial  values  as  compared  to  the  theory. 

The  falling  off  of  the  potential  of  Mixture  A  (normal  hydrochloric  acid 
saturated  with  hydroquinone  and  quinhydrone)  from  the  theoretical 
may  also,  perhaps,  be  attributable  to  reaction  between  the  hydrochloric 
acid  and  the  quinone  (from  the  quinhydrone)  lowering  the  hydrogen-ion 
concentration. 

II.    POTENTIALS  IN  NEUTRAL  AND  ALKALINE  SOLUTIONS. 

As  already  mentioned,  no  definite  results  were  obtained  in  neutral  or 
alkaline  solutions.  The  various  cells  that  were  tried  will  be  taken  up 
individually  for  record  and  to  illustrate  the  general  qualitative  behavior. 
We  will  call  this  series,  Series  IV.  All  of  these  solutions  were  saturated 
with  hydroquinone  and  quinhydrone. 


39 


No.  i. 
(Tenth-normal  Potassium  Chloride.) 


Days. 

Volts.1 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

O 

0-55II 

IO 

0.7136 

2O 

0.7326 

30 

0.7538 

I 

0.5610 

II 

0.7154 

21 

0.7342 

31 

0.7570 

2 

o  .  6640 

12 

0.7170 

22 

0.7369 

32 

3 

0.6615 

13 

0.7183 

23 

0-7379 

33 

0-7547 

4 

0.6727 

14 

0.7214 

24 

0-7399 

34 

5 

0.6828 

15 

0.7240 

25 

0-7435 

35 

0-7547 

6 

0.6932 

16 

0.7244 

26 

0.7438 

36 

0-7545 

7 

0.7070 

17 

0.7257 

27 

0.7461 

37 

0.7541 

8 

0.7088 

18 

0.7282 

28 

0.7501 

38 

0-7543 

9 

0.7114 

19 

0.7304 

29 

0.7523 

On  the  38th  day,  an  apparent  equilibrium  having  been  reached,  the  cell 
was  opened,  while  a  strong  stream  of  nitrogen  was  bubbled  through 
it  to  keep  out  the  air,  and  a  few  drops  of  a  saturated  solution  of  sodium 
hydroxide  were  added.  The  cell  was  closed  again  with  a  moderate  stream 
of  nitrogen  left  running.  Eight  minutes  later  the  potential  had  fallen 
to  0.4194,  and  after  i  hour  it  was  0.4040.  The  nitrogen  was  allowed  to 
run  for  several  hours,  and  then  stopped.  The  table  is  continued  below. 
The  first  column  gives  days  after  start,  and  the  second  column,  days  after 
adding  the  alkali.  In  the  third  column  are  the  potentials. 

i.    n.    in. 

59  21    0.2829 

60  22 

61  23    0.2729 

62  24 

63  25   0.2682 

64  26   0.2712. 

65  27   0.2635 

66  28    0.2607 

67  29    0.2783 

68  30    0.2667 

69  31    0.2659 

Here  again  a  constant  potential  was  finally  reached.  The  contents 
of  the  cell  were,  at  this  point,  titrated  for  alkali,  and  showed  approxi- 
mately 0.034  normal.  Owing  to  the  very  dark  color  of  these  solutions, 
it  is  impossible  to  obtain  an  accurate  titration  of  them.  It  is  necessary 
to  use  litmus  paper,  as  indicator,  and  the  value  obtained  is  only  a  rough 

1  All  potential  values  given  in  this  section  are  single  potential  differences  for  the 
solutions  in  question.  They  are  derived  from  the  observed  electromotive  forces  of  the 
combinations  by  adding  the  value  (0.5265)  of  the  calomel  electrode  when  the  solution 
forms  the  positive  pole,  or  by  subtracting  the  observed  e.  m.  f.  from  0.5625  when  the 
solution  forms  the  negative  pole;  for  in  the  former  case  the  solution  has  a  higher  and  in 
the  latter  a  lower  oxidizing  potential  than  that  of  the  calomel  electrode.  In  which  of 
these  two  ways  a  given  value  was  derived  can  be  determined  by  noting  whether  it  is 
larger  or  smaller  than  0.5265. 


I. 

II. 

III. 

I. 

II. 

ill. 

39 

i 

0.3623 

49 

ii 

0.3243 

40 

2 

0.3661 

50 

12 

0.3174 

4i 

3 

0.3624 

5i 

13 

0.3106 

42 

4 

0.3630 

52 

14 

0.3079 

43 

5 

0.3640 

53 

15 

44 

6 

0-3545 

54 

16 

0.2977 

45 

7 

0.3446 

55 

17 

46 

8 

56 

18 

0.2944 

47 

9 

0-3345 

57 

19 

48 

10 

0.3299 

58 

20 

0.2807 

40 

approximation.  Of  course,  even  an  accurate  titration  would  give  in 
this  case  no  indication  of  the  hydrogen-ion  concentration,  as  the  un- 
limited supply  of  hydroquinone  present  neutralizes  most  of  the  hydroxide 
ion.  Relatively  very  large  quantities  of  hydroquinone  are  required  to 
saturate  these  alkaline  solutions,  where  much  alkali  is  present,  owing  to 
salt  formation  (acid  ionization)  and  resinification.  A  considerable  ex- 
cess of  solid  was  always  added. 

CELL  No.  2. 
(Tenth-normal  Potassium  Chloride.) 


Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

0 

o  .  5409 

II 

0.6993 

22 

0.7117 

33 

0.7420 

44 

I 

0.5571 

12 

0.7016 

23 

34 

0-7439 

45 

0-7431 

2 

0.6287 

13 

0.7027 

24 

0.7199 

35 

o  -  7449 

46 

0-7435 

3 

0.6576 

14 

o  .  7024 

25 

36 

0-7474 

47 

0.7444 

4 

o  .  6707 

15 

0.7049 

26 

37 

5 

0.6786 

16 

0.7071 

27 

38 

0-7450 

6 

o  .  6830 

17 

28 

39 

0.7441 

7 

o  .  6869 

18 

o  .  7084 

29 

0.7305 

40 

0-7445 

8 

0.6915 

19 

0.7097 

30 

0-7335 

4i 

0.7425 

9 

0.6954 

20 

0.7102 

31 

0.7387 

42 

0.7424 

10 

0.6972 

21 

0.7104 

32 

43 

0-7433 

Here  again  a  constant  potential  was  reached  which  was  only  about 
ten  millivolts  lower  than  that  reached  by  Cell  No.  i.  On  the  forty- 
seventh  day  a  few  drops  of  saturated  sodium  hydroxide  solution  we^e 
added,  as  before.  Ten  minutes  later  the  potential  was  0.7424,  after  l/% 
hour  0.7362,  after  3  hours  0.4409,  and  after  6  hours  0.4141.  The 
table  is  continued  below 

i.    ii.    in. 

62  15   0.4731 

63  16    0.4763 

64  17   o . 4808 

65  18    0.4841 

66  19    0.4832 

67  20    0.4864 

68  21    0.4768 

The  contents  of  the  cell  titrated  0.03  normal  alkali.  In  composi- 
tion, therefore,  this  cell  is  almost  a  duplicate  of  No.  i.  Their  potentials 
do  not  differ  greatly  during  the  neutral  stage,  but  ifter  the  addition 
of  alkali  they  do  not  check  at  all.  The  erratic  nature  of  the  alkaline 
potentials  is  further  evidenced  by  the  fact  that,  while  those  of  Cell  No.  i 
decrease  steadily  to  the  constant  potential,  those  of  Cell  No.  2  first  de- 
crease and  then  increase. 

To  see  whether  the  potassium  chloride  exerted  any  influence  upon 
the  potential,  saturated  potassium  chloride  solution  was  tried  instead 
of  the  tenth -normal. 


I. 

II. 

ill. 

I. 

II. 

III. 

48 

•  i 

o  .  4040 

55 

8 

49 

2 

0.3898 

56 

9 

o  .  4405 

50 

3 

0.3892 

57 

10 

0.4465 

5i 

4 

58 

ii 

52 

5 

0.4117 

59 

12 

0-4594 

53 

6 

60 

13 

54 

7 

0.4273 

61 

H 

0.4692 

Ceu,  No.  3. 
(Saturated  Potassium  Chloride.) 

Days.  Volts. 

0  0.4884 

1  0.4880 

2  0.4890 

3  0.4900 

This  potential  being  almost  constant  from  the  start,  though  very 
different  from  the  tenth-normal  KC1  potentials,  three  drops  of  saturated 
sodium  hydroxide  solution  were  added  on  the  third  day.  Twenty-five 
minutes  later  the  potential  was  0.3623  volt.  The  table  is  continued 
below. 

CEU<  No.  3    (Continued). 


I. 

II. 

III. 

I. 

II. 

III. 

I. 

II. 

III. 

4 

i 

0.3346 

ii 

8 

0.2380 

18 

15 

0.2257 

5 

2 

0.2691 

12 

9 

0.2359 

19 

16 

O.2244 

6 

3 

0.2559 

13 

10 

0-2354 

20 

17 

0.2240 

7 

4 

0.2415 

14 

ii 

0.2327 

21 

18 

0.2209 

8 

5 

0.2486 

15 

12 

O.2294 

22 

19 

0.2209 

9 

6 

0.2417 

16 

13 

0.2285 

23 

20 

0.2212 

10 

7 

0.2396 

17 

14 

O.226O 

24 

21 

0.2213 

The.  con  tents  titrated  0.018  normal  alkali.  This  potential,  it  will  be 
noticed,  fell  steadily  to  an  apparent  equilibrium,  which  was,  however, 
lower  than  either  No.  i  or  No.  2,  though  the  cell  contained  less  alkali. 
The  neutral  potential  was  also  lower. 

CELL  No.  4. 
(Saturated  Potassium  Chloride.) 


Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

O 

0-4375 

H 

o  .  5058 

28 

42 

0.5634 

56 

o 

5857 

I 

o  .  4466 

15 

0.5098 

29 

43 

0.5661 

57 

o 

5790 

2 

o  .  4626 

16 

0.5157 

30 

0 

5265 

44 

0.5654 

58 

3 

0.4738 

17 

31 

o 

5533 

45 

o  .  5690 

59 

o 

5913 

4 

0.4776 

18 

0.5237 

32 

46 

60 

5 

0.4874 

19 

0.5261 

33 

0 

5525 

47 

0.5717 

61 

o 

5921 

6 

0.4911 

20 

0.5292 

34 

48 

0-5734 

62 

o 

5919 

7 

o  .  4934 

21 

35 

o 

5587 

49 

0-5733 

63 

0 

5926 

8 

22 

0.5321 

36 

0 

5601 

50 

0.5757 

64 

o 

5899 

9 

23 

0-5347 

37 

o 

5601 

5i 

0.5765 

65 

0 

5827 

10 

0.4965 

24 

0-5349 

38 

o. 

5601 

52 

66 

0. 

5897 

ii 

0.5005 

25 

0.5522 

39 

53 

0.5801 

12 

0.5024 

26 

40 

o. 

5623 

54 

0.5802 

13 

0.5043 

27 

4i 

0. 

5626 

55 

No  alkali  was  added  to  this  cell.  It  is  evident  that  neutral  potentials 
for  this  system  cannot  be  determined  in  potassium  chloride  solutions. 

The  following  alkaline  cells  were  made  up  without  potassium  chloride, 
by  adding  hydroquinone  and  quinhy drone  to  sodium  hydroxide  solutions 


42 

of  various  concentrations,  in  the  electrode  vessel,  while  a  rapid  stream  of 
nitrogen  was  bubbled  through  the  solution.  The  cells  *were  then  closed 
and  the  stream  of  nitrogen  continued  through  the  night  to  facilitate  the 
dissolving  of  the  solids.  The  inlet  and  outlet  tubes  were  then  closed,  as 
with  the  other  cells,  shutting  the  cell  off  from  the  air,  in  an  atmosphere 
of  nitrogen.  Nevertheless  the  solutions  became  dark  immediately  upon 
adding  the  hydroquinone. 


CELL  No.  5. 

Days. 

Volts. 

Days 

.     Volts. 

Days 

.     Volts 

Days. 

Volts. 

Days.       Volts. 

O 

0.2292 

4 

O. 

1495 

8 

O. 

1420 

12 

0. 

1375 

16     0.1335 

I 

0.2315 

5 

0. 

1475 

9 

O. 

1406 

13 

0. 

1342 

17     0.1324 

2 

0.1607 

6 

0. 

1446 

10 

0. 

1405 

14 

0. 

1297 

3 

0.1536 

7 

0. 

1441 

ii 

O  . 

1444 

15 

O. 

1329 

This  cell  titrated  0.4  normal  when  removed,  on  the  iyth  day. 

CELL  No  6. 
(Titrated  i  .6  normal,  when  removed.) 


Days 

.      Volts. 

Days.     Volts. 

Days.       Volts. 

Days.     Volts. 

Days.       Volts. 

O 

o.  1098 

6      0.0711 

12      O  .  0498 

i82   0.0806 

23      O.IO2I 

I 

o.  1464 

7      0.0673 

13      0.0489 

19      0.0740 

24      O.IOI7 

2 

0-I4491 

8     o  .  062  i 

14      0.0475 

-3      O.I5I7 

25      0-.0996 

3 

o  .  0980 

9     0.0612 

15      0.0472 

2O      O.IoSl 

26      0.0971 

4 

0.0850 

10     0.0576 

16     0.0466 

21       O.IO39 

5 

0.0782 

II     0.0564 

17    0.0463 

22       O.I04I 

The  following  might  be  suggested  as  an  explanation  of  the  changes  in  potential 
following  the  stopping  and  resumption  of  the  flow  of  nitrogen  in  the  above  cell.  Un- 
der the  catalytic  effect  of  the  platinum,  hydrogen  is  produced  at  the  electrode,  as  fol- 
lows: 

C6H4(OH)2  —  C6H402  +  2H+  +  20 

2H+  —  H2  +  20 
C6H4(OH)2  ^  C6H402  +  H2 

This  reaction  reaches  equilibrium  when  the  concentration  of  hydrogen  produced  is 
such  that  the  two  systems  represented  by  the  two  partial  or  electrochemical  reactions 
have  the  same  potential.  While  the  nitrogen  is  bubbling  through  the  solution,  at  the 
electrode,  the  hydrogen  cannot  accumulate.  We  have  then  a  resultant  potential,  in- 
termediate between  the  respective  potentials  of  the  two  systems.  When  the  stream 
of  nitrogen  is  stopped,  however,  the  hydrogen  accumulates,  and  the  resultant  potential 
falls  towards  that  of  the  hydroquinone  system.  On  resuming  the  flow  of  nitrogen,  the 
hydrogen  is  again  driven  out  and  the  resultant  potential  rises  again. 

1  The  passage  of  nitrogen  was  stopped  after  taking  this  reading. 

2  The  stop-cock  in  the  bridge  tube,  connecting  the  vessel  with  the  intermediate  KC1 
solution,    slipped,    admitting  air  to  the  bridge  tube.     This  was  forced  out  again  by 
means  of  nitrogen,  without  air  having  been  admitted  to  the  cell,  but  the  cell  contents 
were  disturbed  thereby. 

3  Passage  of  nitrogen  for  three  hours  was  followed  by  a  further  rise  in  potential. 


43 


CELL  No.  7. 
(Hydroquinone  and  Quinone  Added  to  o  .  i 

N  NaOH  to  Saturation.) 

Hours. 

Volts. 

Hours. 

Volts. 

Hours. 

Volts. 

Hours. 

Volts. 

19 

0.3250 

117 

0.2736 

212 

0.2513 

307 

0.2473 

44 

0.3201 

141 

0.2624 

236 

0.2556 

336 

0.2576 

72 

0.3I2I 

170 

0.2601 

26l 

o  .  2460 

365 

0.2544 

93 

0.2914 

190 

0.2581 

284 

0.2493 

38l 

0.2498 

CELL 

No.  8. 

(Duplicate  of  Cell  No. 

7-) 

o 

0.3280 

67 

0.2236 

162 

0.2533 

260 

0.2504 

8 

0.2935 

98 

0.2338 

187 

0.2544 

283 

0.2535 

20 

0.2420 

127 

0.2363 

212 

0.2523 

44 

0.2428 

141 

0.2396 

241 

0.2526 

It  will  be  noticed  that  the  initial  potentials  of  the  above  duplicates 
check  quite  well,  but  that  after  the  first  day  there  is  no  reproducibility. 
In  all  probability,  therefore,  only  the  initial  potential  can  be  taken  as 
representing  the  simple  original  known  system  as  it  was  made  up,  and 
this,  probably  only  approximately;  and  no  definite  significance,  relative 
to  the  original  system  is  to  be  attached  to  any  apparent  equilibrium 
reached  later.  The  later  potentials  are  evidently  more  or  less  accidental. 

The  following  three  cells  were  made  up  similarly  to  the  above,  with 
o.oi  N  sodium  hydroxide  solution: 


Cell  No.  9. 


Cell  No.  10. 


Cell  No.  11. 


Cell  No.  11  (Cont'd). 


Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

0 

0.4027 

O 

o  .  4064 

O 

0.4065 

14 

0.5346 

I 

0.4417 

I 

0.4396 

I 

0-4454 

15 

0.5410 

2 

0.5464 

2 

o  .  4806 

2 

0-4939 

16 

0.5452 

3 

0.4498 

3 

0.4726 

3 

0.4693 

17 

4 

0.4583 

4 

0.4724 

4 

0.4883 

18 

0.5559 

5 

0.4811 

5 

0.5H3 

19 

0-5595 

6 

0.5199 

20 

0.5621 

7 

0.5223 

21 

0-5579 

8 

0.5143 

22 

0.5675 

9 

0.5164 

23 

0.5732 

10 

0.5166 

24 

0.5606 

ii 

0.5176 

25 

0.5598 

12 

0.5265 

26 

o  .  5504 

-' 

13 

0.5265 

27 

0.58II 

It  would  be  interesting  to  compare  these  initial  potentials  wich  the 
van't  Hoff  equation.  But  to  do  this  the  concentrations  must  be  known, 
and  the  determination  of  these  in  alkaline  solutions  is  a  very  different 
proposition  from  what  it  is  in  acid  solutions,  owing  to  the  acid  properties 
of  hydfoquinone,  and  perhaps,  to  some  extent,  of  quinone,  and  their  in- 
stability toward  alkalies.  To  determine  this  data  would  require  a  special 
and  difficult  investigation,  beyond  the  scope  of  this  undertaking. 


44 

Qualitatively  it  is  shown  that  the  oxidizing  potentials  decrease,  that  is, 
the  reducing  potentials  increase  with  alkalinity,  as  would  be  expected 
from  the  theory. 

The  results  in  neutral  and  alkaline  solutions  are  given  largely  as  a 
matter  of  record.  Hardly  more  has  been  done  than  to  give  an  idea  of 
some  of  the  difficulties  confronting  an  attempt  to  study  the  system  quan- 
titatively under  these  conditions.  The  circumstances  here  are  proba- 
bly similar  to  those  encountered  in  the  case  of  the  solutions  containing 
quinone  in  excess  (namely,  disturbing  reactions),  but  to  a  more  pro- 
nounced degree;  and  the  main  disturbing  reactions  are  probably  different. 

One  of  the  difficulties  is  the  determination  of  hydrogen  ion  concentra- 
tion. In  the  case  of  the  acid  solutions  conductivity  data  were  used  for 
this  purpose.  This  involved  the  tentative  assumption  that  the  condition 
of  the  acid,  or  the  hydrogen  ion  concentration,  was  the  same  as  in  a  pure 
solution  of  hydrochloric  acid.  We  cannot  make  a  similar  assumption 
regarding  the  sodium  hydroxide  because  of  the  action  of  the  hydroquinone 
as  an  acid,  toward  alkali.  We  therefore  cannot  use  the  conductivity 
method  in  this  case.  The  question  naturally  arises,  why  not  determine 
the  hydrogen  ion  concentration  by  the  e.  m.  f.  method  with  a  hydroge 
electrode.  This  method  was  not  used  at  all  in  this  work  (although  it  was 
realized  that  the  assumption  upon  which  the  use  of  the  conductivity 
values  was  based  might  be  the  source  of  considerable  error),  because  it 
was  believed  that,  even  if  a  constant  potential  could  be  obtained  with 
a  hydrogen  electrode,  in  the  presence  of  another  active  electrochemical 
system,  this  potential  still  might  be  very  different  from  the  true  hydrogen- 
hydrogen-ion  potential  because  of  the  influence  of  the  other  system.  It 
was  felt  that  this  was  too  big  a  question  to  be  taken  up  merely  as  an  inci- 
dental to  this  investigation. 


PART  III. 

DETAILS   OF   EXPERIMENTAL   METHODS. 

All  volumetric  apparatus  used  was  Eimer  and  Amend  normal  apparatus. 
Each  piece  was  calibrated  and  found  to  be  accurate  to  within  the  limits  of 
precision  possible  in  its  use. 

Standard  hydrochloric  acid  solutions  were  made  up  from  Baker's 
Analyzed  hydrochloric  acid  and  standardized  at  25°  against  Baker's 
Analyzed  sodium  carbonate,  dried  to  constant  weight.  The  manner  in 
which  the  solutions  were  made  up  for  potential  measurement  is  illus- 
trated by  the  case  of  Mixture  A.  Enough  powdered  hydroquinone  and 
quinhydrone  to  just  saturate  100  cc.  of  normal  hydrochloric  acid  was 
weighed  out  and  placed  in  a  100  cc.  volumetric  flask.  50  cc.  of  2  normal 
acid  were  then  added,  from  a  pipette,  at  25°  C.,  and  the  flask  gradually 
diluted  to  the  mark,  keeping  the  contents  at  25°,  with  shaking,  and  later 
whirling,  in  the  constant  temperature  bath,  so  that  when  the  mark  was 
reached  practically  all  of  the  solids  had  dissolved,  and  the  solution  was 
just  normal  to  hydrochloric  acid  and  saturated  with  the  solids  at  25°. 

The  hydroquinone  used  was  partly  Merck's  and  partly  Eimer  and 
Amend's.  Only  one  grade  was  sold  and  it  was  claimed  to  be  very  pure 
because  of  its  method  of  manufacture.  No  accurate  method  of  deter- 
mining its  purity  could  be  found.  It  all  melted  sharply  at  169°  (uncor- 
rected).  This  is  the  melting  point  given  in  Beilstein.1  The  corrected 
melting  point  was  173.0°. 

The  quinone  was  Kahlbaum's.  Some  of  it  was  recrystallized  from 
gasoline  (which  was  found  to  be  an  excellent  solvent  for  separating  qui- 
none from  its  impurities),  and  some  was  sublimed,  both  giving  clean, 
bright  yellow  products,  titrating  the  same  (99.4%  of  the  theory)  and 
melting  sharply  at  115.7°  (corrected)  (the  melting  point  given  by  Beil- 
stein) . 

The  quinhydrone  was  made  in  two  ways;  some  by  treating  hydro- 
quinone with  ferric  chloride,  acidified  with  hydrochloric  acid,  in  aqueous 
solution,  and  some  by  mixing  together  equivalent  quantities  of  hydro- 
quinone and  quinone,  in  aqueous  solution.  The  product,  in  each  case, 
was  filtered  off  and  washed  with  water.  The  product  of  the  first  method 
titrated  99.2%  and  that  of  the  second  99.4%  of  the  theoretical,  the 
latter  checking  the  quinone.  Some  was  recrystallized  from  alcohol  v 
followed  by  ether,  and  some  from  glacial  acetic  acid.  Glacial  acetic  acid 
is  an  excellent  solvent  for  the  purpose,  being  by  far  the  best  of  the  three. 
All  of  the  recrystallized  products  titrated  99 . 4%  of  the  theoretical.  Only 
the  products  giving  this  titration  were  used  in  the  experiments.  Quin- 
hydrone decomposes  on  heating,  before  reaching  its  melting  point. 
1  Hlasiwetz,  A.,  177,  336. 


46 

Nothing  new  or  unusual,  in  the  way  of  apparatus  or  method,  was  in- 
volved in  the  measurement  of  the  potentials.  The  mixtures  whose  poten- 
tials were  to  be  measured  were  placed  in  Kales  calomel  cells.1  The  cell 
and  bridge  tube  were  filled  with  the  solution,  leaving  about  an  inch  and  a 
half  of  free  space  from  the  top  of  the  cell.  Sufficient  mixed  excess  solids 
were  added  to  loosely  fill  the  two  bulbs  at  the  base  of  the  cell,  so  that  the 
electrode  would  be  immersed  in  a  bed  of  solids,  in  order  to  keep  the  solu- 
tion saturated  in  the  region  of  the  electrode.  A  two-hole  rubber  stopper, 
containing  the  electrode  and  a  drawn  out  tube  for  the  introduction  of 
nitrogen,  both  reaching  nearly  to  the  bottom  of  the  cell,  was  inserted.  A 
tapering  gas  outlet  tube,  fitted  with  a  pinch-cock,  was  inserted  in  the  side 
tube  of  the  cell.  Bight  cells  or  less,  at  a  time,  one  of  which  was  the  stand- 
ard cell,  were  grouped  in  a  circle  around  a  beaker  containing  saturated 
potassium  chloride  solution,  their  bridge  tubes  dipping  into  the  solution. 
To  minimize  flow  and  diffusion  the  cock  in  the  bridge  tube  was  kept 
closed  with  a  band  of  vaseline  around  the  upper  and  lower  thirds  and 
the  middle  third  free  for  the  solution  to  pass  around.  The  ends  of  the 
bridge  tubes,  dipping  into  the  salt  solution,  were  plugged  with  wooden 
plugs  wrapped  with  filter  paper.  The  cluster  of  cells  was  immersed  in 
the  constant  temperature  bath.  The  nitrogen  passed  from  the  tank 
through  an  alkaline  pyrogallol  solution,  then  successively  through  water, 
cotton,  a  large  bottle  of  water,  immersed  in  the  constant  temperature 
bath,  acting  as  a  humidifier  (to  prevent  evaporation  in  the  cell)  and  as  a 
manifold,  from  which  the  gas  was  distributed  to  the  various  cells.  The 
closing  of  the  bridge  tubes,  mentioned  above,  did  not  increase  the  resis- 
tance enough  to  interfere  materially  with  the  sensitiveness  of  the  measure- 
ment. 

The  standard  cell  used  was  the  saturated  potassium  chloride  calomel 
cell,  which  was  found  to  be  the  most  satisfactory,  having  by  far  the  great- 
est constancy  and  reproducibility.  It  was  compared  with  a  similar  cell 
belonging  to  Dr.  Fales,  which  had  been  checked  against  a  number  of  other 
standard  cells.  The  comparison  extended  over  a  period  of  one  week. 
The  potential  differences  between  the  two  cells  (our  cell  minus  Dr.  Fales') 
are  given  below : 


Hours. 

Volts. 

Hours. 

Volts. 

Hours. 

Volts. 

Hours. 

Volts. 

o 

0.00042 

27 

0.00033 

72 

O.OOO22 

I24 

0.00034 

I 

0.00042 

31 

0.00030 

74 

O.OOO25 

125 

0.00028 

2 

0.00042 

35 

0.00030 

75 

0.00023 

127 

O.OOO3O 

6 

0.00042 

49 

O.OOO3O 

78 

0.00023 

131 

O.OOOOS 

24 

0.00034 

54 

O.OOO25 

103 

0.00029 

144 

0.00034 

25 

0.00034 

57 

O.OOO29 

I2O 

0.00037 

147 

0.00031 

26 

0.00033 

59 

0.00025 

122 

O.OOO3O 

1  Fales  and  Vosburgh,  J.  A.  C.  S.,  40,  1305  (1918). 


47 

Two  months  later  a  new  standard  cell  was  made  up  and  compared  with 
the  old  one,  which  had  been  in  constant  use  in  the  meanwhile.  The  re- 
sults were  as  follows  (new  cell  minus  old  cell) : 


Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

Days. 

Volts. 

0 

0.00040 

2 

0.00031 

5 

o.ooooo 

8 

0.00016 

'/• 

O.OOOI6 

3 

O.OOO2O 

6 

0.00006 

9 

O.OOOI6 

I 

0.00017 

4 

0.00016 

7 

0.00018 

10 

O.OOOII 

The  measurements  were  made  with  a  Leeds  and  Northrup  potentiom- 
eter, a  sensitive  galvanometer,  and  a  certified  Weston  Standard  Cell. 

A  normal  and  a  Beckman  thermometer  were  used  in  the  constant 
temperature  bath.  They  were  standardized  against  a  Baudin  Normal 
thermometer  (Bur.  Int.  des  Poids  et  Mesures  No.  18,537).  The  bath 
was  electrically  regulated,  and,  as  previously  stated,  the  temperature 
variation  was  less  than  a  hundredth  of  a  degree. 

Details  of  Sohibility  Determinations. 

The  methods  have  already  been  outlined.  Details  of  manipulation, 
etc.,  are  given  below. 

Hydroquinone.  In  the  case  of  water,  as  solvent,  an  excess  of  hydro- 
quinone  was  shaken  repeatedly  with  water,  several  minutes  at  a  time, 
during  a  period  of  several  hours,  in  a  bottle  in  the  constant  temperature 
bath,  and  allowed  to  stand  in  the  bath  over  night.  The  acid  solutions 
were  made  up  as  follows,  so  as  to  be  at  the  desired  normality  after  satura- 
tion with  hydroquinone :  About  5  grams  of  hydroquinone  were  placed 
in  a  100  cc.  volumetric  flask,  50  cc.  of  acid  of  twice  the  desired  normality 
were  added  and  thoroughly  shaken,  and  then  enough  water  was  added 
with  shaking  to  dissolve  all  of  the  hydroquinone.  Small  portions  of 
powdered  hydroquinone  were  then  added,  alternately  with  the  water, 
all  at  25°,  so  that  the  solution  was  kept  saturated  with  hydroquinone 
as  it  was  diluted  to  the  mark.  It  was  then  transferred  to  a  bottle,  with 
more  hydroquinone,  shaken  intermittently,  as  in  the  previous  case,  and 
allowed  to  stand  over  night,  or  long  enough  to  leave  a  clear  supernatant 
solution.  Five  cc.  of  the  solution  were  then  pipetted  off  into  a  weighed 
10  cc.  distilling  flask,  fitted  with  a  small  rubber  stopper  and  fine  drawn 
out  tube,  extending  to  the  bottom  of  the  flask,  as  is  customary  in  vacuum 
evaporations,  to  aid  ebullition.  It  was  then  connected  with  a  vacuum 
pump  and  evaporated  to  constant  weight,  at  room  temperature  or  with 
gentle  warming  (not  over  35°),  at  a  pressure  of  about  8  mm.  Heating 
much  over  40°  at  this  pressure  may  cause  appreciable  evaporation  of  the 
hydroquinone.  The  fact  that  concordant  results  were  obtained  from 
different  solutions,  shaken  different  lengths  of  time,  was  taken  to  indi- 
cate that  the  solutions  were  saturated. 


48 

No  suitable  analytical  method  for  the  determination  of  hydroquinone 
could  be  found,  but  it  is  possible  that  the  reduction  of  silver  ion  in  neu- 
tral nitrate  solution  might  be  developed  into  such  a  method  when  the 
proper  conditions  are  worked  out.  In  ammoniacal  solution,  quinone 
will  also  reduce  silver. 

Quinone  and  quinhydrone. — The  way  in  which  the  solutions  were 
saturated  and  sampled  has  already  been  completely  described  under  the 
section  on  Solubilities,  etc.  Only  details  of  the  analytical  method  remain 
to  be  given. 

Solutions. — Tenth-normal  potassium  dichromate  made  from  pure,  dried 
potassium  dichromate,  to  be  used  as  the  standard.  10%  colorless  potas- 
sium iodide  solution.  Tenth,  hundredth  and  five -hundredth  normal  sodium 
thio sulphate  solutions  standardized  against  the  dichromate  solution. 
The  end  point  is  very  delicate. 

Method  of  Standardization  or  Analysis. — To  the  sample  of  solution 
to  be  titrated,  a  number  of  cc.  of  10%  KI  solution,  approximately  equal 
to  the  number  of  cc.  of  Af/io  thiosulphate  that  would  be  required,  were 
added.  Half  this  number  of  cc.  of  concentrated  hydrochloric  acid  were  then 
added  and  the  solution  titrated  immediately,  using  starch  as  indicator,  if  de- 
sired. Usually  the  end  point  is  sufficiently  delicate  without  it.  In  some 
cases  the  exact  starch  end  point  is  rendered  uncertain  by  the  pinkish  color 
probably  caused  by  the  action  of  the  hydrochloric  acid  on  the  quinone. 

Except  for  the  last  mentioned  difficulty,  which  only  causes  trouble  in 
the  case  of  the  very  dilute  solutions  (quinhydrone  in  HC1  solutions  with 
considerable  excess  hydroquinone)  and  is  not  very  serious  then,  the  method 
seems  to  be  a  very  accurate  one.  Repeated  determinations  on  purified 
quinone  and  quinhydrone,  varying  the  concentrations  somewhat,  with- 
out going  to  extremes,  always  gave  the  same  result,  viz.,  99.3-99.5% 
of  the  theoretical.  The  remaining  0.6%  may  be  moisture  or  a  constant 
error  in  the  method.  No  statement  or  data  as  to  whether  the  method  is 
supposed  to  give  theoretical  results  or  not  could  be  found  in  the  litera- 
ture on  the  subject. 

CONCLUSION. 

It  has  been  found  that  whenever  the  concentrations  could  be  con- 
trolled and  side  reactions  minimized  (namely,  in  the  acid  solutions  with 
small  proportions  of  quinone)  fairly  constant  and  reproducible  potentials 
were  obtained  with  solutions  containing  hydroquinone  and  quinone,  in 
equilibrium  with  their  addition  compound,  quinhydrone.  Very  likely, 
with  elaborate  and  specially  designed  apparatus  and  carefully  studied 
precautions,  still  more  definite  results  could  have  been  obtained.  The 
present  work  being  more  or  less  of  a  pioneer  nature,  great  precision,  in 
this  respect,  was  not  attempted,  the  necessary  conditions  not  having  yet 
been  worked  out. 


49 

These  potentials  were  found  to  vary  with  the  probable  approximate 
concentrations  of  the  substances  involved  in  the  electrochemical  equa- 
tion, in  fair  accordance  with  the  van't  Hoff  constant  temperature  equa- 
tion. 

In  alkaline  solutions  no  quantitative  study  was  possible,  owing  to  in- 
terfering reactions,  but,  qualitatively,  the  oxidizing  potential  was  found 
to  be  much  lower  than  in  acid  solutions,  and  to  fall  with  increasing  alka- 
linity in  accord  with  the  theory.  Potentials  in  neutral  solutions  were  in- 
termediate between  those  in  acid  and  in  alkaline  solutions,  but  were  in- 
definite and  greatly  affected  by  the  electrolyte. 

Although  no  case  could  be  found  better  adapted  to  such  an  investiga- 
tion the  conditions  are  nevertheless  such  that  the  concentrations  of  the 
substances  involved,  including  hydrogen  ion,  could  only  be  approximated. 

To  this  end,  the  solubilities  of  hydroquinone,  quinone  and  quinhydrone, 
and  the  probable  approximate  dissociation  constants  of  quinhydrone,  in 
water  and  in  dilute  hydrochloric  acid  solutions,  have  been  determined. 
Part  of  this  has  merely  been  a  more  thorough  repetition  and  extension  of 
work  already  done  by  others.  The  new  fact,  however,  was  brought  out, 
that  the  solubility  product  law,  when  applied  to  saturated  quinhydrone 
solutions,  in  the  presence  of  an  excess  of  hydroquinone,  fails  to  hold  for 
concentrations  of  hydroquinone  approaching  saturation,  owing  evidently 
to  an  increase  in  the  dissociation  constant,  which  becomes  so  abrupt  at 
the  saturation  point  that  the  law  is  actually  reversed,  that  is,  that  addi- 
tion of  hydroquinone  causes  an  increase,  instead  of  a  decrease,  in  the 
quantity  of  quinhydrone  which  will  go  into  solution.  This  cannot  be 
attributed  to  error  in  the  determination  of  the  solubility  of  hydroquinone, 
since  the  change,  though  sudden,  is  a  smoothly  accelerated  change  with  a 
gradual  beginning.  The  abruptness  of  the  change,  however,  suggests  a 
peculiar  condition  of  what  Washburn  calls  "the  thermodynamic  environ- 
ment,"1 connected  with  the  fact  of  saturation.  It  was  also  found  that 
hydrochloric  acid  decreased  the  solubility  of  hydroquinone  but  increased 
that  of  quinone. 

The  confirmation  of  the  van't  Hoff  equation  in  a  case  of  this  kind  per- 
forms two  distinct  functions:  (i)  It  verifies  the  Law  of  Mass  Action  for 
the  electrochemical  equation,  thus  furnishing  evidence  supporting  the 
use  of  the  electrochemical  equation.  (2)  It  affords  a  means  of  evaluating 
the  intrinsic  oxidizing  tendency  (represented  by  the  constant  term  of  the 
van't  Hoff  equation),  which  may  be  called,  for  convenience,  "the  normal 
potential,"  (i.  e.,  the  potential  for  unit  concentrations  or  concentration 
ratio  of  1)  and  which  forms  a  basis  for  comparing  the  system  with  other 
systems,  independently  of  concentrations,  and  for  calculating  the  potential 
for  any  given  concentrations. 

1  Principles  of  Physical  Chemistry. 


50 

Since  the  concentrations  were  only  more  or  less  approximated,  and  since 
the  holding  of  the  ideal  gas  laws  in  the  case  of  osmotic  pressure,  upon  which 
the  van't  Hoff  equation  is  based,  is  also  only  an  approximation,  very  close 
agreement  with  the  theory  could  not  be  expected.  As  it  was,  the  agree- 
ment was  of  the  same  order  as  that  obtained  for  parallel  work  in  inorganic 
chemistry. 

Thus  the  quantitative  relationships,  upon  which  the  electrochemical 
theory  of  oxidation  is  based,  have  been  extended  experimentally  into  the 
field  of  organic  chemistry,  and  in  this  way  the  chemistry  of  organic 
non-electrolytes,  has,  in  the  only  instance  so  far  investigated,  been  corre- 
lated with  that  of  inorganic  electrolytes. 


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